Recycling of alum used for phosphate removal in domestic wastewater treatment


Material Information

Recycling of alum used for phosphate removal in domestic wastewater treatment
Physical Description:
xii, 211 leaves : ill. ; 28 cm.
Cornwell, David Alan, 1948-
Publication Date:


Subjects / Keywords:
Coagulation   ( lcsh )
Water -- Pollution   ( lcsh )
bibliography   ( marcgt )
theses   ( marcgt )
non-fiction   ( marcgt )


Thesis--University of Florida.
Includes bibliographical references (leaves 206-210).
Statement of Responsibility:
by David Alan Cornwell.
General Note:
General Note:

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Source Institution:
University of Florida
Rights Management:
All applicable rights reserved by the source institution and holding location.
Resource Identifier:
aleph - 000159295
notis - AAS5620
oclc - 02624531
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Full Text







to my wife and son


Appreciation is expressed to my committee chairman, John Zoltek,

Jr., for the help which he provided in overcoming the many obstacles in

the path to obtaining this degree. With great admiration, I thank Prof-

fessor T. deS. Furman who sets a standard of excellence for all engineers

to strive for. I am grateful for the guidance and stimulation rendered

by J. Edward Singley and the other member of my committee, D. 0. Shah.

I will never forget the encouragement given during difficult times

nor the enjoyment shared during successful times by J.S.T.

Had it not been for the guidance of my parents I might never have

had the striving necessary to obtain this degree.

Words cannot express my gratitude to my wife, Linda, who sacrificed

so that I could obtain this degree. It would not be possible for some-

one to have been more understanding and encouraging during these times

than Linda has been.

This research was supported by a grant through the Engineering and

Industrial Experiment Station funded by the State of Florida.








1-1. Sources and Quantities of Phosphorus 1
1-2. Standards for Phosphorus Removal 2
1-3. Phosphorus Removal 3
1-4. Purpose of This Research 5

2-1. Chemistry of Aluminum Phosphate Precipitation 6
2-2. Phosphorus Removal by Aluminum Addition to Secondary
Effluent 17
2-3. Aluminum-Phosphate Sludge 17
2-4. Aluminum-Phosphate Sludge Dissolution 19
2-5. Recovery of Alum in Water Treatment Plants 23
2-6. Recovery of Alum in Wastewater Treatment Plants 29
2-7. Rational for Current Research 33

3-1. General Description 34
3-2. Classification of Extraction Processes 35
3-2-1. Organic separations 36
3-2-2. Inorganic Separations 37
3-3. Metal Extraction by Alkyl Phosphoric Acids 39
3-4. Theoretical Approach to Metal Extraction 43
3-4-1. Equilibria of metal extraction 43
3-4-2. Kinetics of metal extraction 45
3-4-3. Mechanism of metal extraction 46
3-5. Engineering Approach to Metal Extraction 50
3-5-1. Equilibria qf metal extraction 50
3-5-2. Kinetics of metal extraction 55
3-6. Pertinent Solvent Extraction Literature 57

4. L:f ER 1 i El; ; APPARATUS AND t '.':iCDi!.E i
4-1. Feed Solution
4-1-1. Synthetic feed solution
4-1-2. Tertiary wastewater sludge feed solution
4-2. Alkyl Phosphoric Acids
4-3. Analytical Equipment and Techniques
4-3-1. Total phosphate determination
4-3-2. Orthophosphate determination
4-3-3. pH determination
4-3-4. Aluminum determination
4-3-5. Available H+ determination
4-3-6. Emulsion size determination
4-3-7. Total solids determination
4-3-8. Volatile solids determination
4-4. Experimental Procedure and Equipment
4-4-1. Treatment of wastewater sludge samples

5-1. Aluminum Extraction by Di(2-ethylhexyl) Phosphoric
5-2. Aluminum Extraction by an Equal Molar Solution of
Mono- and Di(2-ethylhexyl) Phosphoric Acid
5-3. Mechanism of Aluminum Extraction by an Equal Molar
Solution of Mono- and Di(2-ethylhexyl) Phosphoric
5-3-1. Existance of the alkyl phosphate at the
5-3-2. Chemical reaction at the interface
5-3-3. Entrance of the metal complex into the
kerosene phase
5-3-4. Summary of proposed mechanism for aluminum
extraction by MDEHPA
5-4. Stripping of Aluminum-Mono-Di(2-ethylhexyl) Phos-
phoric Acid

6-1. Acidification of the Aluminum-Phosphate-Organic
6-2. Extraction of Aluminum from the Acidified Sludge
6-3. Stripping of Aluminum from the MDEHPA Solution
6-3-1. Stripping at low aluminum concentrations
6-3-2. Stripping at high aluminum concentrations
6-3-3. Summary of aluminum stripping

7-1. Full-Scale Operation at a 60 mg/Z Alum Dose
7-1-1. Sludge pre-treatment
7-1-2. Sludge acidification
7-1-3. Aluminum extraction
7-1-4. Disposal of aqueous phosphate solution
7,-1-5. Stripping of the extract
7-1-6. S.iFar' of alum recovery at a 60 mg/i alum















PA r E

7-2. Full-Scale Operation at 200 a-./. Alum dose 187
7-2-1. Sludge pre-treatment 187
7-2-2. Sludge acidification 187
7-2-3. Aluminum extraction 187
7-2-4. Disposal of aqueous solution 188
7-2-5. Stripping of the extract 188
7-2-6. Summary of alum recovery at a 200 mg/i
alum dose 191

8-1. Case I. 60 mg/Z Alum Dose Without Alum Recovery 192
8-2. Case II. 60 mg/L Alum Dose With Alum Recovery 194
8-2-1. Chemical costs at 60 mg/Z alum dose 194
8-2-2. Total capital and operating costs at
60 mg/i alum dose 195
8-3. Case III. 200 mg/i Alum Dose Without Alum Recovery 197
8-4. Case IV. 200 mg/i Alum Dose With Alum Recovery 197
8-4-1. Chemical costs at 200 mg/i alum dose 197
8-5. Summary of Process Costs 198

9-1. Summary of Alum Recovery by Solvent Extraction 201
9-2. Theoretical Findings of This Research 203
9-3. Suggested Further Research 203
9-3-1. Theoretical considerations 203
9-3-2. Engineering considerations 204





2-1 Properties of Commercial Alum 6

2-2 Aluminum to Phosphate Molar Ratio :Eceary for Phosphate
Removal 11

2-3 Aluminum and Phosphate Equilibrium Constants 15

2-4 Coagulation of Raw Water by Filtrate from Acid-Treated
Sludge Palin's Data 24

2-5 Aluminum Recovery as a Function of pH Webster's Data 27

3-1 Abbreviations of Acid Organophosphorus Compounds 42

4-1 Characteristics of Secondary Effluent of the University
of Florida Campus Sewage Treatment Plant 69

4-2 Physical and Chemical Properties of Alkyl Acid Phosphates 71

4-3 Solubility Properties of Alkyl Acid Phosphates 72

5-1 Acceleration of DEHPA Extraction of Iron (III) by Proton-
Accepting Complexing Agents 93

5-2 MDEHPA Extraction of Aluminum Dimer Reaction Range 114

5-3 MDEHPA Extraction of Aluminum Monomer Reaction Range 115

5-4 Stripping of 0.1 M MDEHPA Containing 0.023 M Aluminum 143

6-1 Acid Requirements for Dissolution of Aluminum 149

6-2 Aluminum Recovery as a Function of pH 156

6-3 Extraction of Aluminum From Acidified Aluminum-Phosphate
Sludge 157

8-1 Cost of Chemicals Used in Aluminum Recovery by Solvent
Extraction 193

8-2 Aluminum Recovery by Solvent Extraction Capital and
Operating Costs 196

8-3 Summary Cost of Aluminum Recovery by Solvent Extraction 199



2-1 Aluminum-Phosphate Solubility Diagram at an Ionic
Strength of Zero 14

2-2 Aluminum-Phosphate Solubility Diagram 21

3-1 Schematic Diagram of Alkyl Phosphoric Acids 41

3-2 Diagram of Metal Extraction as Proposed by the
Two-Film Theory- 49

3-3 Graphical Construction for Solvent Extraction Design
with Immiscible Liquids 53

4-1 Batch Simulation of Three-Stage Continuous
Countercurrent Cascade 65

4-2 Phosphate Removal as a Function of Alum Dose 68

5-1 Proposed Molecular Structure for a) DEHPA Dimer, b) 1:4
UO2-DEHPA Complex, c) Polymeric U02-DEHPA Complex 80

5-2 Percent of Tributyl Phosphate (TBP) Needed to Prevent
Third Phase Formation of Calcium-DEHPA (DEHPA-Ca) Complex 85

5-3 Percent Aluminum Extracted by DEHPA as a Function of pH 87

5-4 Extraction Coefficient of Aluminum as a Function of DEHPA 89

5-5 Equilibrium Curve for Aluminum Extraction by DEHPA 92

5-6 Effect of Contact Time on Extraction of Aluminum by
0.1 M MDEHPA 97

5-7 Extraction Coefficient of Aluminum by !DEHiPA as a
Function of pH 100

5-8 Equilibrium Curve for Aluminum Extraction by 0.1 M MDEHPA 102

5-9 Equilibrium Curve for Aluminum Extraction by 0.2 M MDEHPA 104

5-10 Equilibrium Curve for Aluminum Extraction by 0.3 M MDEHPA 106

5-11 Equilibrium Curve for Aluminum Extraction by 0.4 M MDEHPA 108



5-12 MDEHPA Necessary to Extract Greater than 96% of the
Initial Aluminum in the Feed 111

5-13 Titration Curve of 0.12 M !MEHPA 118

5-14 Droplet Size Frequency Distribution for Kerosene
Sprayed at 700 psi 124

5-15 Hydrogen Bonding for Monoalkyl Phosphoric Acids 127

5-16 Schematic Diagram of Proposed Mechanism of Aluminum
Extraction by MDEHPA 131

5-17 Effect of Cetyl Trimethyl Ammonium Bromide (CTAB) on
Critical Micelle Concentration (CMC) of Potassium
Laurate (KL) 133

6-1 Titration Curve of Aluminum Phosphate Sludge with
Sulfuric Acid 148

6-2 Scehmatic Representation of the Three Major Zones
of Sludge Settling 152

6-3 Settling Curves of Aluminum-Phosphate Sludge After
pH Reduction 154

6-4 Effect of Contact Time on Percent Aluminum Stripped from
MDEHPA at Low Aluminum Concentrations 161

6-5 Effect of Contact Time on Percent Aluminum Stripped from
MDEHPA at High Aluminum Concentrations 164

6-6 Pseudo-Equilibrium Curve for Stripping of Aluminum
from MDEHPA with 6 N H2S04 167

6-7 Solubility Diagram for Hydroxy Apitite and Dicalcium
Phosphate 169

6-8 Phosphate Remaining in Solution Using Lime as pH Control
and as the Source of Calcium Ions 171

7-1 Flow Diagram of Alum Recovery Process by Solvent Extraction 177

7-2 Graphical Determination of Number of Stages Necessary for
98% Aluminum Extraction at an Initial Alum Dose to
Secondary Effluent of 60 mg/i 181

7-3 Graphical Determination of Number of Stages Necessary for
97% Aluminum Stripping at an Initial Alum Dose to
Secondary Effluent of 60 mg/ 185




7-4 Graphical Determinaticn of Number of Stages Necessary for
98% Aluminum Stripping at an Initial Alum Dose to
Secondary Effluent of 200 mg/Z 190

Abstract of Dissertation Presented to the
Graduate Council of the University of Florida
in Partial Fulfillment of the Requirements for the
Degree of Doctor of Philosophy



David Alan Cornwell

August, 1975

Chairman: John Zoltek, Jr.
Major Department: Environmental Engineering Sciences

A procedure was developed for the economical recovery of aluminum

when the aluminum was used as a coagulant for phosphorus removal in

domestic wastewater treatment. The process was developed initially

using synthetic solutions. This allowed large quantities of sludge to

be readily available and solution characteristics were easily changed.

Sewage treatment plant sludge was then used to predict the prototype

operating parameters. Aluminum-phosphate-organic sludge was collected

by coagulating secondary effluent from the University of Florida Campus

Sewage Treatment Plant. The sludge produced was found to be a combina-

tion of sterrettite and organic solids.

The sludge was first thickened using a slow-stirring apparatus to

obtain a solids concentration of about four times that of the raw sludge.

The sludge was then reacted with sulfuric acid to dissolve the aluminum

and phosphate. In order to reach a pH of 2.0, two moles of H2SO4 per

mole Al were used. If the supernatant was separated from the residual

organic sludge by sedimentation, 93% of the available aluminum was sepa-

rated. )if vacuum filtration was used, 1002 of the aluminum was separated

from the organic sludge.

In order to separate the acidified aluminum from the phosphate a

solvent extraction process was developed. A kerosene solution of alkyl

phosphates was contacted with the aluminum-phosphate solution in a mixer.

The alkyl phosphates reacted with the aluminum, causing the aluminum to

become kerosene soluble. Several different alkyl phosphates were studied

in order to optimize aluminum extraction. An equal molar mixture of

mono- and di(2-ethylhexyl) phosphoric acid was found to be the most effi-

cient extractant. It was possible to extract 98% of the aluminum from

the acidified aluminum-phosphate solution. At optimal aluminum extrac-

tion conditions the reaction was found to follow an equilibrium equation

such that the alkyl phosphate reacted as a dimer. The value of the

equilibrium constant was 46.40.95 (pK = -1.670.01).

The kerosene and water were very insoluble and separated readily in

a settler. The aluminum rich kerosene phase was contacted with 6 N H2SO4.

The aluminum transferred from the kerosene phase and entered the acid

phase. The kerosene:acid volume ratio was adjusted to produce a final

aluminum concentration equal to the aluminum concentration in commercial

alum (about 5%). Approximately 98% of the aluminum in the kerosene was

transferred into the acid phase in two countercurrent stages. The recov-

ered aluminum was reused as a coagulant for phosphorus precipitation.

The aluminum free kerosene was recycled to the extraction stages.

The overall aluminum recovery was between 89% and 93%. The complete

process was estimated to require tank capacities equivalent to ten 55-gal-

lon drums per million gallons of plant flow. Very little capital invest-

ment would be required and the process would be equally applicable to both

large and small treatment plants. For a treatment plant using 200 mg/l alum

for phosphate removal, the cost of sludge handling with the recovery process

was reduced to 20% of the cost of sludge handling without the recovery process.




Phosphorus nitrogen and carbon are the primary nutrients responsible

for eutrophication of lakes and streams. Wastewater treatment plants

throughout the country are being faced with the need to upgrade their

treatment in order to remove 90, of the biochemical oxygen demand (BOD)

and suspended solids which are indicators of the carbon content of

the effluent stream. In addition to carbon removal, phosphorus and

nitrogen removal is required by many states. The nutrient removal

process easiest to control and the most efficient nutrient removal system

is phosphorus precipitation by chemical coagulation. Coagulation for

phosphate precipitation usually upgrades the carbon removal such that

957, BOD and suspended solids removal standards can be met. All states

that have streams tributary to the Great Lakes have established

:astewater effluent phosphorus limits or require a fixed percentage

reduction of phosphorus. Several other states have also established

phosphorus standards.

1-1 Sources and Quantities of Phosphorus

Domestic wastewater normally has a substantial conc-ettration of

phosphorus. Hum. ratese, including food disposal, account for about

30 f)'} of the phosphorus present in domestic wastcwatersc [1J. Detergents

contri ning pho. p1-phate builders which are used principal ly for lau i2dar-ing

clothes account for the remainder of the phosphorus. Sources of

phosphorus other than the home that may cause deviation from these

percentages include: 1) sodium hexametaphosphate used in corrosion

control of water distribution system's, 2) fertilizer, 3) animal

feedlots, 4) commercial laundry wastes, 5) dairy wastes, and

6) slaughter house wastes. The amount of phosphorus produced by

industrial sources varies considerably from location to location.

The quantity of phosphorus resulting from human excretions ranges

from 0.5 to 2.3 Ib per capital per year [2], with a mean value of 1.2 lb

per capital. The mean annual contribution of phosphorus from synthetic

detergents is about 2.3 lb per capital [3]. The phosphorus values in

domestic wastewater generally range from 6 mg/ to 20 mg/l, with an

average of 10 mg/ [4].

1-2 Standards for Phoisphorus Removal

As of June 1971 [3], sixteen states had adopted wastewater effluent

phosphorus standards. Either an effluent concentration limit, 0.1 mg/f

to 2 mg/-P, or a specified percentage reduction, 80% to 95%, has been

established by these standards. Specifying a final effluent concentration

would be the preferred method for control of eutrophication, since the

effluent concentration determines the concentration in the receiving

water, which in turn governs algae growth.

Dryden and Stern [5] conducted a study to determine the phosphate

concentration needed in the wastewater plant's effluent to prevent

algae growth of the recreational lake into which the effluent was

discharged. The initial phosphorus concentration of 40 mg/ol was reduced

by coagulation with 300 mg/, of alum to less than 0.5 mg/t-P, averaging

about 0.25 mg/L-P. Effluent phosphate concentrations of 0.25 mg/ to

0.3 mg/f-P were diluted upon entry into the lake. The lake phosphate

concentration was less than 0.05 mg/Z-P. Laboratory studies showed that

phosphate concentrations below 0.5 mg/f-P inhibited algae growth to

some extent, but concentrations of 0.1 mg/, to 0.3 mg/.-P significantly

inhibited the rate of growth and the density of the algae. Growth almost

completely stopped at less than 0.05 mg/t-P. It was observed that

between 1 mg/Z and 30 :T /f-P the same algae growth occurred.

Ideally, standards would be established to limit the effluent

phosphate concentration to that necessary to limit the receiving streams'

effluent concentration to a value below that of about 0.05 mg/e.

In some lakes and streams 100%-P removal of wastewater effluents would

not reduce the phosphate concentration below 0.05 mg/e due to runoff

sources of pollution. In other areas, the dilution factor may be high

enough that no treatment plant phosphate removal is required.

1-3 Phosphorus Removal

Phosphorus can be removed from wastewater effluents by ion

exchange, reverse osmosis and other demineralization techniques, or

by chemical coagulation. The most economical method is chemical

coagulation. The three major precipitants for phosphorus removal

include lime and salts of iron and aluminum. The coagulant can be

added before the primary sedimentation tank, in the aerator of an

activated sludge pia'it, before the final sedimentation tank, or in a

tertiary process. PhiY oshorus removal a~d secondary y effluent clcarifica t on

have been found to be better with alum than wii-h lime [I,6]. Alum and

iron appear to produce about the same phosphorus and suspended solids

removal [7]. The sludge produced from coagulation by either iron or

alum is of very low density, contains large amounts of water of

hydration, and is therefore difficult to dewater. Alum sludge is

generally easier to dewater than iron sludge and consequently alum

is preferred for coagulation [7]. On the other hand, lime sludge dewaters

quite readily and can be disposed of very economically. Recovery

processes have also been developed for lime, lowering the sludge handling

costs further. Lime has therefore often been the coagulant of choice for

phosphate precipitation, even though the effluent quality produced by

lime coagulation is often below that which can be obtained by alum


It has been estimated [8] that a solids content for alum sludge of

20% is needed to landfill the sludge in conjunction with other materials

such as municipal solid wastes, and that a 40% solids content is necessary

for landfilling the sludge alone. Present methods of alum sludge

dewatering require large capital and operating expenses. The cost of

alum sludge handling can be twice as much as the cost for capital,

operation, and maintenance of the coagulation process for phosphate

precipitation by alum added to secondary effluent.

The use of alum as a tertiary coagulant would probably increase

if an economical method of sludge disposal could be found. Our nation's

recent interest in the three R's, recovery, recycling, and reuse, has

led many investigators to the conclusion that some form of alum

recovery and recycling scheme is needed for alum treatment to become

economically justified. There have been no economical systems developed

to recover alum from sludges where phosphate removal is one of the

treatment plant objectives. The aluminum can be dissolved easily

and separated from the organic sludge, but the phosphorus is also

dissolved. Before the aluminum can be reused as a coagulant, it must

be separated from the phosphate. The high cost of aluminum-phosphate

separation has prevented alum recovery from becoming economical.

1-4 Purpose of This Research

It was the purpose of this research to devise an economical means

of recovering alum from aluminum phosphate sludges produced by coagulation

of secondary-treated wastewater. The recovery process requires that

the alum be phosphate free, and in a form suitable for reuse as a

coagulant. The alum should be concentrated sufficiently to minimize

storage and pumping costs. The process should be adaptable to small

treatment plants at a minimal initial capital outlay. Finally, the

process should be easy to operate, and compatible with existing waste-

water treatment processes.

Solvent extraction has been successfully used in the recovery

of some heavy metals from leachate liquors. In this research a solvent

extraction recovery process has been developed for aluminum recovery

from sludges produced by aluminum addition to secondary effluent for

phosphate precipitation.

Chapter 2 describes the present state of the art of alum recovery,

while Chapter 3 discusses solvent extraction as applicable to this

research. The procedures that were employed are outlined in Chapter 4.

The process was first applied to synthetic sludge solution, Chapter 5,

and then to actual wastewater treatment plant sludges, Chapter 6. The

proposed process design is presented in Chapter 7, followed by an economic

evaluation in Chapter 8.



In this chapter the chemistry of aluminum phosphate compounds are

presented, including the reactions pertinent to aluminum phosphate

coagulation and subsequent sludge dissolution. Also presented is a

brief review of operating data for phosphorus removal by alum.

Characteristics of aluminum phosphate sludge are discussed. Finally

a literature review of previous methods to recover alum is presented.

2-1 Chemistry of Aluminum Phosphate Precipitation

The most abundant forms of dissolved inorganic phosphorus exist
as protonated orthophosphates, H PO protonated pyrophosphates,

H P2-( protonated trimetaphosphates, H P 0(3x), and protonated
HxP207 x 3 9
tripolyphosphates, HxP3010 In secondary sewage effluents the

protonated orthophosphates represent 70 to 90% of the total phosphorus

present [9]. Theoretically all polyphosphates in the natural environment

eventually hydrolyze into orthophosphate. Consequently the interaction

between the soluble orthophosphate group and the cation used for

coagulation is of particular importance in phosphate removal from

wastewater effluents.

Orthophosphate falls under the general category of a polyprotic

acid. With a tribasic acid such as phosphoric acid, there are three

stages of ionization. The d .:e of ionization is dependent upon the

pH of the solution. In the pH range from 6 to 8, the ringe encountered

in secondary :*:.*-, the phosphate is present primarily as H2PO4 and


The principal aluminum compounds that are commercially available

and suitable for phosphorus precipitation are alum and sodium aluminate.

Both of these chemicals are available in either liquid or dry forms.

Alum is acidic in nature while sodium aluminate is alkaline. By far

the most widely used aluminum compound for coagulation is alum. Almost

all the alum produced in the United States is manufactured from bauxite

and bauxitic clays. The ore is ground and reacted with sulfuric acid.

The reaction is represented by the following equation,

Al203(s) + 3H2S04 + 11H20 = Al2(SO4) 314110 (2-1)

The liquor is separated from any unreacted ore and then adjusted in

strength to 8% to 8.3% Al203. It is practically free of insoluble natter

and sold as liquid alum. The purified alum solution can be converted

to dry alum by carefully controlled evaporation so that the Al203 content

is brought up to 17%. The solution is spread on long slabs and allowed

to solidify. The resulting solid is then ground or powdered and shipped

in bags or in bulk. Typical properties of liquid and dry alum are given

in Table 2-1.

Liquid alum is shipped at about a 50% alum concentration. Liquid

alum can be fed directly to the coagulation basin or can be diluted

into the range of 15% to 50% before use as a coagulant feed. It has

been shown [10] that dilution has little effect on the efficiency of

coagulation by alum. As pointed out by Recht and Ghassemi [11], dilution

of the coagulant feed may not affect the coagulation properties of

TA3LE 2-1

Properties of Commercial Alum

Chemical formula

Molecular weight

% A1203

Weight/gal, lbs/gal

Dry alum equivalent, Ibs/gal

Bulk density, lbs/cu ft

pH of 1% solution

Solubility in water
At 680F
At 320C

Crystallization point, OF


Al2(S04) 314H20




Al2 (S04) 3 14H20







87 gms Al2(SO4)3-14H20/100 gms H20
71 gms Al2(S04)3'14H20/100 gms H20


Source: American Cyanamid Co., Cyanamid Alum, p. 40, 1972.

aluminum due to the fact that the aluminum feed solution does not

undergo hydrolysis below a pH of 3.0. However, economics play an

important factor in the choice of concentration of the alum solution

to be fed. For example, if a plant was feeding an average of 250 mg/Z

alum from a 0.1% alum solution, the required alum flow rate would be

25% of the plant capacity. In order to minimize pumping costs and

the resulting plant capacity a concentrated solution of alum is highly


Despite considerable research, the basic chemistry of phosphate

reactions with Al3+ has remained obscure, and interpretations of data

reported in the literature have often been quite contradictory. To

illustrate, Lea et al. [7] and Henriksen [12] presented results

supporting the view that the removal of phosphate involved adsorption

onto precipitating aluminum hydroxide. However, according to Stumm [13],

Cole and Jackson [14] and Recht and Ghassemi [11], the interaction of

aluminum with orthophosphate resulted in the formation of insoluble

metal phosphates. Considerable disagreement also exists between various

investigators on the kinetics and stoichiometry of aluminum-phosphate

reactions, and on the effect of various parameters such as pH and ionic

concentrations on the efficiency of phosphate removal. Thus, Stumm [13]

reported that under proper pH control and at low aluminum-phosphate

ratios, the reaction of aluminum with orthophosphate was

A13+ + H2P04- = AlP04(s) + 2H+ (2-2)

provided that sufficient time for the precipitation was allowed. In

actual practice, however, even under optimum pH conditions, considerably

higher than stoichiometric amounts of aluminum are needed for complete

precipitation of phosphates. In order to account for the higher

aluminum necessary for phosphate precipitation, the following reactions

have been proposed to occur, depending upon the alkalinity of the


Al2(SO 4)314H20 + 6HCO3 = 2A1(OH)3(s) + 6CO2 + 14H20 + 3SO42- (2-3)


Al2(SO )314H O = 2Al(OH)3(s) + 3H2 SO + 8H20 (2-4)

Farrell et al. [15] presented the following general equation for

phosphate precipitation by alum.

A12(S04)3 + Na2HPO4 + 3H20 = Al(OH)3*AlPO4(s) + Na2SO4 + 2H2S04 (2-5)

Equation (2-5) gives an aluminum to phosphorus ratio of 2:1. Approximate

molar ratios necessary for phosphate precipitation are presented in

Table 2-2 [3]. Cole and Jackson [14] analyzed the sludge resulting

from aluminum phosphate precipitation. By x-ray differaction they found

two forms of aluminum phosphate: variscite, Al(OH)2H2P04; and sterrettite,

[A1(OH)2]3HPO4H2PO4. From their results the following reactions can be


3A12(SO4)3*14H20 + 2H2PO4 + 2HPO4 = 2[Al(OH)2]3HPO4H2PO (s) +

6H2SO4 + 3SO 4 + 30H20 (2-6)


A12(SO4)3-14H20 + 2H2PO4 = 2A1(OH)2H2PO4 (s) + 2H2S04 + SO2-+10H20 (2-7)

Reaction (2-6) has an Al:P ratio of 3:2 and in reaction (2-7) the ratio


Aluminum to Phosphate Molar Ratio
Necessary for Phosphate Removal

P Reduction Al:P Alum:P
Required Mole Ratio Weight Ratio Weight Ratio

75% 1.38:1 1.2:1 13:1

85% 1.72:1 1.5:1 16:1

95% 2.3:1 2.0:1 22:1

Source: EPA, "Process Design Manual for Phosphorus Removal," contract
#14-12-936, Oct. 1971, p. 3-3.

is 1:1. A log solubility diagram at an ionic strength of zero (I=0) is

shown in Figure 2-1 for AlPO4(s) and Al(OH)2112PO4(s). Table 2-3 shows

the pertinent equilibrium constants. An equilibrium constant was not

available for sterrettite and consequently a solubility diagram could

not be drawn for this compound. As can be seen variscite has a lower

solubility than AlP04, with only a slight difference in minimum

solubility pH. Recht and Ghassemi [11] found an Al:P molar ratio of

1.4:1 to be optimum for phosphate removal and concluded that one of the

ionized aluminum hydroxide forms reacted with the phosphates rather

than free Al3+. Their work showed evidence that reaction 2-6) may

have been controlling the precipitation. Cole and Jackson [14] found

Al(OH)2H2PO4 to dominate when coagulation took place at a pH of 2.4

to 4.0; whereas [Al(OH)2]3HPO4H2P04 dominated at a coagulation pH of

5.0 to 6.0, which is the optimal pH range for phosphate removal in

wastewater effluents. Undoubtedly the precipitation is more complicated

than that represented by reaction (2-6), but this reaction may be more

indicative of the compounds formed than are reactions (2-2) through

(2-5). A possible source of aluminum demand other than carbonates and

phosphates is the reaction of aluminum and organic anions. Struthers

and Sieling [16] found that hydroxyl groups such as tartrate and oxalate

and carboxylic groups such as citrate and malate inhibited precipitation

of aluminum phosphates. Before precipitation of aluminum phosphates

took place the aluminum first had to react with all the organic anions.

Many organic alcohols and acids are by-products of enzyme action and

would be expected to be present in wastewater effluents.

The optimum pH for aluminum-phosphate coagulation has been reported

to be between pH 5.5 and 6.5 [3]. It can be seen from Figure 2-1

that the optimal pH for AlPO4 and Al("'1)2HL' 4 precipitation is about 6.0.

Figure 2-1.

Aluminum-Phosphate Solubility Diagram at an Ionic
Strength of Zero. Diagram drawn such that the
aluminum:phosphate molar ratio in solution equals
1:1. The species considered were A1PO4(s),
Al (OH) 2"H;04 (s), A10H2+, and A1(OH)4.






15 .
j O*r .iinnli niJ

2 3 4 5 6 7 8 9 10 11 12 13


Aluminum and ?:c;ophate Equilibrium Constants

Ionic Strength
Species I = 0 I = 1

AlPO4(s) 10-21 10-18.3

Al(OI)2H2P04(s) 10-30.5 10-28.7
Al3+ + 4(OH)- = A1(OH)4 1032.5 1030.7

3+ = UO12i
Al3+ 20 = A1012+ 10-5 10-5.6

H3P4, Kal 10-2.12 10-2.42

Ka2 10-721 1i-7.81

Ka3 10-12.32 10-13.22

Rccht and Ghassemi [11] conducted kinetic experiments on the

precipitation of orthophosphate with aluminum nitrate. The data in-

dicated a drop in phosphate concentration from the initial 12 mg/Z-P

to 0.1 mg/-P in less than 60 seconds following the addition of aluminum

nitrate. No further removal of phosphate was observed after this

period, despite the very noticeable gradual growth and agglomeration of

the precipitate. They concluded that the reaction itself was very rapid

and that the kinetics of aluminum-phosphate precipitation were dependent

upon mixing.

The effectiveness of aluminum salts as coagulants for polyphosphate

removal has also been the subject of controversy. Sawyer [17] reported

that alum salts are highly effective in removing all forms of phosphates.

According to experiments reported by Stumm [13] tripolyphosphates were

not removed to an appreciable extent due to the formation of soluble
complexes such as A1P30 Recht and Ghassemi [11] found the optimal

pH range for aluminum-polyphosphate coagulation to be very narrow, with

the optimum at pH 5.5. Even at a 2:1 A1:P ratio complete polyphosphate

precipitation was not achieved. Tripolyphosphates could only be reduced

by 83% under optimum conditions. Pyrophosphates were reduced by 95%

at a pH of 5.5. Pyrophosphate removal also required an Al:P ratio of

2:1, which may account for Al:P ratios greater than 1:1 being necessary

for phosphate removal from secondary effluent.

Sodium aluminate has met with mixed success as a coagulant for

phosphate removal.. The sodium a.luminate-phosphate reaction can be

represented as follows,


NaA102 + PO4

+ 2H 0 APO
2 APO(s) + NaOH + 30H1


with the accompanying reaction,

NaA102 + H2CO3 + H20 = Al(OH)3(s) + NaHCO3 (2-9)

Unlike alum, sodium aluminate raises the pH upon addition. Low

alkalinity waters do not respond well to phosphate removal by sodium

aluminate, because the pH rises above the optimal pH range for phosphate

removal [15].

2-2 Phosphorus Removal by Aluminum Addition to Secondary Effluent

Equipment for phosphate removal typically consists of mixers,

flocculators, settlers and filters. An alum dosage of 100 to 200 mg/C

is required for phosphate removal from most wastewater effluents,while

dosages of 50 to 100 mg/ are generally sufficient for effluent clari-

fication of suspended solids. Settling alone will reduce residual

phosphate to 1 mg/, while filtration can reduce phosphates to less

than 0.1 mg/. Surface overflow rates for clarifiers have ranged from

580 to 1440 gpd/ft with the lower end of this range giving more

consistent results. Filtration rates of 2 to 5 gpm/ft2 have been used.

For a more detailed analysis of phosphorus removal facilities the reader

is referred to the"EPA 1971 Process Manual"[3].

2-3 Aluminum-Phosphate Sludge

The sludge flow produced by aluminum coagulation of secondary

wastewater effluents is approximately 1% of the total plant flow [18].

The sludge has a solids content of about 0.5% to 1%. If a 100 mg/e

alum dosage is used and aluminum loss in the effluent is considered

negligible, the sludge will contain an equivalent of 10,000 mg/e alum,

of 910 mg/t Al 3+. Typically volatile organic solids represent about

40% [19] to 50% (this research) by weight of the total sludge. Aluminum-

phosphate sludge can be expected to increase the total plant sludge

flow by about 50% [20] by volume and double the dry weight solids [19,20]

produced by the treatment plant.

Aluminum-phosphate sludges have presented significant disposal

problems. The primary methods for disposal have been digestion or

dewatering followed by landfilling. It has been reported [3] that sludges

resulting from aluminum addition during secondary treatment could be

used in anaerobic digestion. However, Brown [21] found that the alum

sludge had an adverse effect on the anaerobic digestion process and that

the resulting digested sludge would not dewater well on drying beds.

Attempts at several wastewater plants in Florida [20] to use aluminum

coagulated sludges in conjunction with primary sludge in digestors

were not successful. The result was that the plants were forced to use

aerobic digestion. Aerobic digestion has been successful in producing

a 5% solids content sludge [20]. However, a large increase in digestion

facilities was necessary to accommodate the increase in mixed liquor

suspended solids.

Dewatering waste alum sludge is difficult. The vacuum filter yield

of alum wastewater sludge is 0.4 to 0.8 lb/hr/sq ft as compared to 10

to 40 lb/hr/sq ft for lime wastewater sludge and 4 to 8 lb/hr/sq ft for

mixed digested sludge. A weight solids content of 20% for alum sludge

is near the maximum that can be attained by vacuum filtration.

Several othermethods are available for alum sludge dewatering.

Most of the processes have been used in the water treatment industry

rather than wastewater. Several methods of alum sludge dewatering

were compared by Westerhoff and Daly [22]. Pressure filtration, scroll

and basket centrifugation, and artificial freeze-th;nw were all e -:loyed.

Pressure filtration had operating characteristics very similar to vacuum

filtration, with a 40% final solids content being possible. Pressure

filtration was shown to be less expensive for dewatering alum slu'-.-

than vacuum filtration [22]. Scroll centrifugation produced a 20% final

solids and basket centrifugation produced a final solids of about 10%.

Neither centrifugation method has found widespread use for alum sludge

dewatering. A freeze-thaw process was used to improve sludge dewater-

ability. In this process the sludge was frozen, thawed, and thickened.

A 17% final solids was attained by this method. If the thickened sludge

was vacuum filtered a 40% to 60% solids material was produced. The

freeze-thaw method was not found to be economical. Presently vacuum

filtration is the most common method of alum sludge dewatering, with

pressure filtration gaining acceptance as a viable alternative.

2-4 Aluninul-Ph._p1,.tL. Sludge Dissolution

Aluminum-phosphate solid is amphoteric and can be dissolved either

in acid or alkaline. Figure 2-2 shows the theoretical AlPO4 solubility

diagram at a ionic strength of 1. If AlPO4(s) is the primary solid

present, at a pH of 2 approximately 5x10-1 M solution of Al3+ can be

dissolved in equilibrium with AlP04, whereas a 5x10-1 M solution

cannot be dissolved on the basic side until a pH of 12.6 is reached.

It is less expensive to lower the pH to 2.0 than to raise the pH to

12.6 and therefore the acidic method for alum dissolution has generally

been utilized in water treatment. The dissolution can be represented by,

2A1(OH) (s) + 3H SO = Al (SO ) + 6H 0
3 2 4 2 4 3 2


Figure 2-2.

Aluminum-Phosphate Solubility Diagram at an Ionic
Strength of One. Diagram drawn such that the
aluminum:phosphate molar ratio in solution equals
1:1. The species considered were AlP04(s),
AI(OH)2H2?04(s), A1OH+, and Al(OH) .


when phosphates are not present, and by

2AlPO4(s) + 3H2S04 = A12(SO4)3 + 2H3PO4 (2-11)

when phosphates are present. In the field of wastewater treatment

both the acidic and basic methods of AlPO4 dissolution have been

attempted. Generally, the alkali dissolution can be represented by

reaction (2-12) in potable water treatment,

Al(OH)3(s) + NaOH = NaA102 + 2H 0 (2-12)

and by reaction (2-13) in wastewater treatment.

AlPO (s) + 4NaOH = NaAlO2 + Na3PO + 2H20 (2-13)

If phosphate removal is an objective of the treatment it is necessary

to separate the dissolved aluminum, NaA102 or A12(SO4)3, from the

dissolved phosphate, Na3PO4 or H PO before the aluminum is reused

as a coagulant. The majority of the research on alum recovery has been

done in the field of potable water treatment where alum was used for

turbidity and color removal, and phosphates were not a porblem. Alum

recovery has found application in water treatment plants in Japan,

but the United States has been slow in adopting its use. The major

problem with recovery in the water treatment field has been the redissol-

ving and therefore recirculating of the precipitated color and the

possible build-up of inorganic impurities. The next section will

present a review of the research that has been done on alum recovery

in potable water treatment. This will be followed by a review of the

alum recovery studies conducted on wastewater alum sludges where phosphate

removal was a treatment objective.

2-5 Recovery of Alum in Water Treatment Plants

As outlined by Roberts and Roddy [23] the earliest attempt to

reclaim alum sludge was made by Jewel, who in 1903 patented a process

for water treatment and for reclaiming the coagulant by reacting the

aluminum hydroxide with sulfuric acid. Mathis, in 1923, was issued

a patent for basically the same process as developed by Jewel. "Black

Laboratories," of Orlando, Florida, in 1951 :1cested the use of an

alum sludge recovery process utilizing the sulfur dioxide gas from boiler

stacks as a source of sulfuric acid. Some of the first reported alum

recovery research in the water treatment field was by Palin' [24].

Palin's work was conducted at the Whittle Dene Waterworks of Newcastle,

England. In his first set of experiments filter washwater was treated

with 0.05% and 0.1% (by volume) sulfuric acid. Chloride was added to

oxidize the color present in the washwater. This treated washwater

was then used in conjunction with fresh alum in order to determine

the amount of fresh alum needed to lower the color to 10 Hazen. Palin

found that in treating raw water 28 ppm alum was needed, while treating

raw water plus 3% treated washwater only 11 ppm alum was needed.

Despite this large reduction in alum dosage the cost of acid used was

higher than the cost of alum saved. In a second experiment, 0.4% of

concentrated sulfuric acid was added to sludge with a 2% solids content.

The results of the subsequent coagulation are shown in Table 2-4.

Palin reported superior results when the sludge was charred at 400 C

before acid treatment. It was found that 1 ton of oven dried sludge

would yield 2 tons of aluminum sulfate cake (14% Al203) upon addition of

about 0.9 tons of 98% H2SO4'
2 4


Coagulation of Raw Water by Filtrate from
Acid-Treated Sludge Palin's Data

Percent treated sludge
added to raw water 0.0 0.13 0.25 0.38

Color, Hazen 44.0 47.0 27.0 0.0

pH 7.5 7.2 6.9 6.8

Source: Palin, A. T., Proc. Soc. Wat. Treatment Exam, 3:131 (1954).

In Tampa, Roberts and Roddy [23] studied the recovery of alum in

both a pilot and full scale process. The recovery was based upon the

following reaction,

2A(OH) 3(s) + 3H2SO = A12(SO )3 + 6H20 (2-14)

The alum sludge was thickened by settling for 3 hours. The solids

content reached 1% in pilot plant scale and as high as 2% on full

scale operation. The sludge samples were reacted with enough sulfuric

acid to convert the aluminum hydroxide to aluminum sulfate. The amount

of acid used varied depending upon the alkalinity of the rax water.

The pH range for complete aluminum dissolution was between 1.5 and 2.5

for highly alkaline and less alkaline waters, respectively. After the

reclaimed alum was recycled ten times, there was no reduction in

finished water quality. It was estimated that chemical costs could be

reduced by 70% using the acid recovery method.

Isaac and Vahidi [25] in 1961 studied alum recovery for a method

of sludge disposal, noting the difficult problem of drying and disposing

of the alum sludge. Isaac first tested the alkaline and acid methods

of aluminum recovery. He found that aluminum recovery with caustic soda

is never very satisfactory. It was also observed that organic matter

bound with the aluminum hydroxide, especially organic color, was much more

soluble in alkali than in acid. It was therefore decided to use the

acidic method for aluminum recovery. Aluminum was recovered from fresh

sludge and from anaerobically digested sludge. Tests were then conducted

to determine the volume occupied by the sludge after acid treatment.

At a pH of 2.5, corresponding to 79% aluminum recovery, a 74% volume

reduction of sludge was obtained. Recovered alum was usually about

75% as efficient as fresh alum in reducing color, although one test

resulted in an efficiency of 89%. The researchers concluded that the

pH should be lowered to about 3.0 for a recovery of about 60Z to 65%

of the aluminum since the organic color was not dissolved to an ex-

cessive extent at this pH.

In laboratory experiments Webster [26] found that if sulfuric acid

were added to alum sludge to depress the pH value to about 2.4, a

clustering effect of the floc particles took place with extremely rapid

settling of the insoluble matter. The supernatant liquor contained the

alum, representing about 80% recovery. Table 2-5 shows aluminum recovery

versus pH. A pilot plant for alum recovery was then constructed. Good

coagulation was not obtained with recycled alum that had been recovered

at a pH of below 3.0. Webster concluded that the alum reduced the pH

of the raw water below the range for acceptable color removal. Therefore,

the pH of the sludge was reduced to 3.5 for alum recovery and reuse.

No detrimental effects resulted from continued recycling of the alum

recovered at a pH of 3.5.

Fujita [27] noted that the new water treatment plant at Asaka, Japan,

used an alum recovery system. Part of the sludge handling included

acid treatment followed by thickening with alum recovery. The remaining

sludge was sent to vacuum filtration. No operating data were presented.

Streicher [28] conducted pilot tests to determine the usefulness

of acid recovery of aluminum followed by filter pressing the remaining

organic sludge. The pH was reduced to 1.5 to 2.5 by sulfuric acid.

He found that when the ratio of Al(OH)3(s) to other suspended matter

in the sludge was high, considerably less than stoichiometric amounts of

sulfuric acid were required. If the ratio were low, more than


Aluminum Recovery as a Function of

pH Webster's Data

Sample Sludge ml 25% pH % Alum
Volume, ml H2SO4 Recovery

I 1000 3.5 3.77 53.7

II 1000 4.0 3.63 64.7

III 1000 4.5 3.50 70.7

IV 1000 5.0 3.21 78.6

V 1000 5.5 2.98 78.2

VI 1000 6.0 2.64 81.3

Source: Webster, J. A., J. Inst. Wat. Eng. 20:2, 167 (1966).

stoichiometric amounts of acid were needed. Acid treatment resulted

in reduction of sludge volume to less than 10% of the original volume,

and a concentration of the sludge to 20% solids. The alum recovery

was 80% to 93%. With the use of a filter press the remaining organic

sludge was concentrated to 40% to 50% solids.

Fulton [29,30] described an alum recovery system scheduled to be

put into operation in 1974 at Jersey City, New Jersey. The process

consisted of thickening, acid addition and filter pressing of the

resulting sludge. The acid recovery could be bypassed and only the

filter press used if necessary. For a 100 MGD plant, the savings were

estimated at $4.60 per million gallons when alum recovery was used.

An alum recovery of 90% was estimated.

Westerhoff [31] in 1973 conducted a 15-week pilot plant study to

determine the effect of recycling alum recovered from waste alum sludge

by an acidic process. The pH of the sludge was reduced to 2.0 for

conversion of aluminum hydroxide to aluminum sulfate. The main purpose

of the study was to evaluate potential contaminant build-up in the

recycled alum. Measurements were made on total microscopic count,

coliform, hardness, alkalinity, cyanide, fluoride, phenol, dissolved

solids, nitrates, sulfates, chlorides and several metals such as copper,

lead and zinc. Throughout the study final water analysis for the pilot

plant using recycled alum and for the full scale plant using fresh alum

were essentially the same, indicating that impurities were not built-up

by recycling alum..

Westerhoff and Daly [8,22,32] conducted a complete study of various

alum sludge dewatering facilities. They included pressure filtration

with and without alum recovery, centrifugation, rotary vacuum filtration,

horizontal vacuum filtration with and uitho'ut alum recovery, coaiguiation,

filter press and freeze-thaw. The studies showed alum recovery followed

by horizontal vacuum filtration to be a workable process warranting

economic evaluation. Th.- recovery of alum varied from 50% to 90%.

Coagulation-basin sludge was thickened from an initial 4% to 6% solids

to a final 21% solids by acid treatment. After filtration the solids

content was 37%. However, because of the low alum dosage used for raw

water water turbidity removal, the most economical method of alum

sludge treatment was determined to be pressure filtration without alum


2-6 Recovery of Alum in Wastewater Treatment Plants

When treating domestic wastewater with alum,one of the major

objectives is the reduction of phosphates. When the resulting sludge

is treated with either acid or base to dissolve the aluminum, the

phosphate is also returned to solution. Therefore, before the aluminum

can be recycled the phosphate must be removed from the alum solution.

When the acidic method of aluminum dissolution has been used, the

phosphate has been removed by ion-exchange or absorption. When the

aluminum was dissolved by the addition of alkali, the phosphate was

removed by precipitation with calcium.

The first reported attempt to recover alum from wastewater sludges

was made by Lea [7] in 1954. Lea used the alkaline recovery method

resulting in the formation of sodium aluminate as governed by the

following reactions,

Al(OH) -[PO 4 (s) + 4HaOH = NA102 + Na PO4 + 2H2 + 30H (2-15)
3 4 2 3 4 2

NaAl2 + 2Na3PO4 + 3CaCl2 = NaAO12 + Ca3(PO )2(s) + 6NaC1 (2-16)

Raising the pH to 11.9 brought 93% of the aluminum into solution. After

calcium chloride addition for precipitation of tricalcium phosphate

only 12% of the phosphates remained in solution. The reaction for reuse

of sodium aluminate was

NaAlO2 + H2CO + H2 = Al(OH) (s) + NaHCO3 (2-17)

The wastewater being treated in this study contained sufficient carbon

dioxide to hydrolyze the sodium aluminate without further acid addition.

Advantage could be taken of the fact that make-up alum applied directly

to the raw water will supply additional acidity, as shown by the

following equation:

Al2(SO ) -14H20 = 2A1(OH) (s) + 3H2SO (2-18)

Unfiltered effluent had 77% to 89% phosphate removal and filtered samples

had 93% to 97% phosphate removal using recycled sodium aluminate as

the coagulant. It was estimated that chemical costs could be reduced

by 60% to 90% depending upon whether the precipitated calcium phosphate

sludge was marketable as a fertilizer.

In 1966 Slechta and Culp [6] attempted several methods of aluminum

recovery followed by phosphate removal from the recovered alum solution.

Recovery was first done with sodium hydroxide and calcium chloride, as

in Lea's work. While Lea recovered 93% of the alum at pH of 11.9,

Slechta and Culp achieved only 85% recovery after filtration. Slechta

found that after addition of calcium chloride to the recovered alum

solution, the final aluminum concentrations were substantially reduced

when the pH was between 10 and 11.5. It was also found that occasionally,

after addition of calcium chloride, the entire solution would turn into

a white, gelatinous mixture, which did not dewater easily by sedimentation

or filtration. This occurred in the pH range 10 to 12, and it appeared

that a calcium aluminate was precipitating. This process was not found

to be economical. Alum recovery was next attempted using lime in place

of sodium hydroxide. The maximum recovery occurred at a pH of 10.1.

Aluminum recovery by sedimentation was 19.1%, and by filtration, 35.2%.

This method was also found to be uneconomical. Contrary to.findings

by Lea, Slechta and Culp did not find sodium aluminate to be a good

coagulant. The equivalent alum dose had to be increased to 400 mg/

using 50% alum and 50% sodium aluminate to obtain the same residual

turbidity as obtained with 100% alum at a dose of 200 mg/f. i.l--i 30%

alum and 70% sodium aluminate were used, good turbidity removal could

not be obtained up to an equivalent alum dose of700 mg/-. The alkaline

method of alum recovery was abandoned and an acid method was attempted.

Essentially 100% alum recovery was possible at a pH of 2, and from 95%

to 100% recovery was possible at pH 2.5. Two methods were studied to

remove phosphates from the reclaimed alum: ion exchange and activated

alumina. With the anionic ion exchange resin it was found that 53.5

tons of regeneration salt were needed per ton of reclaimed alum,

rendering this scheme uneconomical. Activated alumina with a surface

area of 300 m2/g was used to rempve phosphate from a feed solution

containing 250 mg/Z aluminum and 440 mg/ phosphate. After the effluent

phosphate concentration reached 100 mg/ the alumina was regenerated.

Several methods of regeneration were attempted, including heating and

acid washing. It was not possible to reactivate the alumina to an

extent that acceptable phosphate removal was achieved. It was concluded

that at present there are no suitable methods for alum recovery from

wastewater treatment plant sludges when phosphate removal was one of the

treatment plant objectives.

Farrell et al. [15] in 1968 used a slightly different approach to

the alkaline method of aluminum recovery. The first test used calcium

hydroxide as in the work by Slechta. Maximum aluminum recovery was 45%

at a pH of 10.5. A mixture of sodium hydroxide and calcium hydroxide

was next used for alum recovery. The overall chemical reaction was

reported to be,

3A1(OH)3'AIPO4(s) + 6NaOH + 5Ca(OH) = 6Na[Al(OH)4] + Ca5(PO 4)3OH(s)


Maximum recovery occurred at a calcium hydroxide/aluminum molar ratio of

1.0 and a sodium hydroxide/aluminum molar ratio of 2.5. At these ratios

recovery was 82% at a pH of approximately 11.0. Farrell found that the

recovery was limited by a calcium/aluminum precipitate similar to that

encountered by Slechta and Culp. From ion product calculations it was

suggested that this compound was of the form Ca2 Al'OH. The authors

proposed that the reason Slechta and Culp achieved poor coagulation with

alum-sodium aluminate mixtures was because solutions high in aluminate

produced pH values above the optimum. They noted that for soft waste-

waters there is a narrow pH range for effective phosphate removal. No

attempt was made to use sodium aluminate as a coagulant in the work by


2-7 Rationale for Current Research

From reviewing the literature it was concluded:

1. The alkaline methods did not give high aluminum recovery

values. Sodium aluminate required additional acid addition

for the coagulation of some secondary effluents. The acid

method of recovery therefore appeared most desirable.

2. Conventional methods of aluminum phosphate separation were

not economically feasible.

3. For ease of storage, minimization of pumping costs and plant

size the alum should be concentrated before being recycled.

4. The i-rreainin solids need to be concentrated to about 20% for

disposal. Some disposal areas may require 40%.



3-1 General Description

Liquid extraction, sometimes called solvent extraction, is the

separation of the constituents of a liquid solution by contact with

another insoluble liquid. If the substance distributes itself differently

between the two liquids then separation can take place, and this can

be enhanced by the use of multiple contacts. For example, if a solution

of water containing an organic acid such as acetic acid is contacted

by agitation with ethyl acetate, some of the acid, but very little of

the water, will enter the ester phase. The ester and water have

different densities so that upon settling they may be decanted from each

other. By this method separation of the acid has been achieved. The

remainingwater may be repeatedly contacted with more ester to reduce

the acid content further.

In all such operations the solution which contains the solute to

be extracted is called the feed, and the liquid with which the feed is

contacted is called the solvent. The solute rich solvent product is

called the extract, and the residual liquid from which the solute has

been removed is called the raffinate. It is often desired to extract

one substance from the feed solution while leaving others behind, hence

the term selective extraction. Some processes may use two solvents to

separate the cumponeaLs of one feed solution. For example, a mixture

of p and o nitrobenzoate may be separated by distributing them between

the insoluble liquids chloroform and water. The chloroform pre-

ferentially dissolves the para-isomer and the water the ortho-isomer.

This is called double solvent or fractional extraction.

After extraction the solute must generally be removed from the

solvent. This is often done by distillation or by a second extraction

operation. The removal of solute from the extract by a second extraction

is called stripping. From an economic standpoint the solvent must be

recycled with a minimum of loss.

3-2 Classification of Extraction Processes

Two broad categories of extraction systems can be distinguished

depending on the origins of the differential solubility. In the first,

the differential arises from purely physical differences such as polarity.

Many separations in the organic field belong to this category. As would

be expected of a process depending only on the physical effects of

differences in molecular structure, such separations are rarely very

specific and the separation factors obtained are usually only modest.

In the second category, the differential solubility is due to one of

the solutes interacting chemically with the solvent to form a complex.

This is exploited for many metallurgical separations.

The distinction between the two categories is significant in

influencing both the physical changes which occur during extraction and

the method of solvent recovery. In systems where the differential

solubility arises purely from physical factors, the relative miscibility

of the two phases will usually be a function of the solute concentration.

On the other hand, when definite compound formation is involved, the

mutual solubilities of the two solvents often do not vary significantly

and so the flow rates of the solvents, calculated on a solute free

basis, do not change in passing through an extraction. This constant

solubility simplifies both design calculations and the mathematical

representation of the overall process, which can be important for

considerations of control.

When an extraction process depends upon chemical interaction,

subsequent solute-solvent separation demands reversal of the complex

forming reaction, usually by chemical means. In the case of systems

exploiting only physical forces, solute-solvent separation is achieved

by physical means, usually distillation. Thus the final solute-solvent

separation stage may introduce either a chemical or a thermal cost item

into the overall economics of a process.

3-2-1 Organic separations

The organic chemical industry used the earliest large-scale

applications of solvent extraction in the coal-tar field and for separation

of aromatics from aliphatics in the petroleum industry. The use of

solvent extraction in the organic field does not appear to have expanded

as much as in the inorganic field. However, if measured in terms of

tonnages processed, organic extractions far exceed inorganic.

While there are important exceptions, many separations in the

organic field rely on physical rather than chemical interaction between

the solute and solvent. The solvent extraction processes which have

reached industrial application in the organic field are of a wide

variety. Some of the'application include: 1) petroleum and petrochemical

industries, e.g.,separation of aromatics and paraffins for the production

of aromatics, mercaptan removal from petroleum fractions and the production

of acetic and acrylic acid; 2) the coal-tar industry, e.g., the

extraction of phenols from coal-tar distillates and the separation

of phenol homologues; 3) the pharmaceutical industry, e.g., penicillin

production; 4) soap production; 5) pyrethrunrecovery; and 6) lactic

acid manufacture.

3-2-2 Inorganic separations

The development of the nuclear industry gave tremendous impetus

to research on complex formation between metal salts and organic solvents.

Such solvents can be divided into three broad groups according to whether

their complex formation reaction is one of cation exchange (as with

carboxylic acid), anion exchange (as with amines), or adduct formation

(typified by ethers and neutral organophosphorus compounds).

In adduct formation neutral oxygen-bearing organic solvents are

considered to extract electrically neutral species by virtue of solvation.

The two general types of extractants in adduct formation are carbon-

bonded and phosphorus-bonded oxygen bearing solvents. In organophos-

phorus ester systems, water is often eliminated from the organic phase

metal complex, whereas in ethers and ketones water is a necessary part

of the complex, usually acting as a bond between the solvating molecules

and the inorganic salt. A commonly used process is in uranium extraction

by tributyl phosphate (TBP) represented by the equation (3-1) [33].

UO2(NO3) (aq) + 2R PO(org) = U02(NO3 2(R3PO)2(org) (3-1)

The solvents dealt with here have been used primarily for the extraction

of metal halides and halo-metallic complexes. Although TBP has received

most of the attention as a neutral solvent, a wide variety of other

compounds have been used, including dialkyl diethers, ketones, alcohols,

phosphonates, phosphinates, and phosphine oxides.

Solvent extraction by long chain aliphatic amines has become a

very popular specific field of research. The consensus is that the

chemistry of extraction by amine salts is comparable to the "absorption"

of metal complexes on anion-exchange resins [34]. The extractability

of metals by amines depends more on the aqueous phase conditions than

on the specific affinities of anionic metal complexes towards the type

of the bulky alkylammonium cation. Species which have been extracted

include UO2(NO3)3 FeC14, HC12, and many similar metal-salt combinations.

The extractive power of amines increases from primary amines to secondary

amines to tertiary amines. There are several potential applications of

amine extraction including fuel reprocessing, hydrometallurgy, and

analytical chemistry of metals and in equilibrium studies of metal


Reactions involving cation exchange include carboxylic and phos-

phoric acid extractants. Generally carboxylic acids have relatively

high aqueous solubility and low extractive power. In spite of these

limitations they have been useful in separations of metals having similar

chemical properties by adjustment of the aqueous phase pH. Carboxylic

acids extract a large number of metals primarily from alkali solutions.

Metal extraction exhibits cation-exchange properties of the type

MIX + MX = M2X + MIX (3-2)

where the superscript bar represents the organic phase. Metal extraction

by orgophosphorus compounds has received a great deal of attention.

A number of mono- and dialkyl phosphoric, phosphonic, and phosphinic

acids are available for extraction. Uranium has been the metal (most

studied for extraction by phosphoric acids, and several full scale

operations are in existence. However, several other metals have been

extracted including the lanthanides, iron, chromium, zirconium, titanium,

thorium, and plutonium. The extraction of metals by alkyl phosphoric

acids is of such importance to this research that the next section is

devoted to a general description of this area.

3-3 Metal Extraction by Alkyl Phosphoric Acids

Acidic organophosphorus extractants have proved to be very useful

in the extraction of a wide variety of metal species. The range of

usefulness of these extractants is from purely laboratory procedures

(analysis, separation of radioactive tracers, etc.) to large-scale

processes (recovery of uranium from liquors, isolation of transplutonides

and fission products from highly radioactive wastes,etc.).

Typical molecules of acidic organophosphates are shown in Figure 3-1.

The alkyl groups may or may not be identical in the case of dialkyl

phosphoric acids. One or both of the alkyl groups may be aryl, aroxyl,

alkyl, or alkoxyl, or a mixed alkylaryl, alkylaroxyl, arylalkyl or

arylalkoxyl. A list of common organophosphoric compounds is in Table 3-1.

The functional group, =P(O)OH, is common for all members of this class

of extractants. Due to the presence of both electron donor and electron

acceptor groups in the =P(O)OH grouping, it is typical of all acidic

organophosphates to undergo various specific interactions like self-

association and molecular complex'formation with diluent or other solutes.

As with most systems involving mass transfer and/or chemical

reaction two important operating considerations are the kinetics and

Figure 3-1. Schematic Diagram of Alkyl Phosphoric Acids.

a. Mono(2-ethylhexyl) phosphoric acid

b. Di(2-ethylhexyl) phosphoric acid

C 2

O =P-O-H







0 0
i I


Abbreviations of Acid Organophosphorus Compounds

Abbreviations of acidic organophosphorus compounds are based on simple
regularities: D denotes di, M denotes mono, and primes denote groups
bound directly to the phosphorus without an oxygen atom,

DAPA di-n-amyl phosphoric acid
DBPA di-n-butyl phosphoric acid
DBzPA dibenzyl phosphoric acid
DClPhPA di(p-chlorophenyl) phosphoric acid
DEHPA di(2-ethylhexyl) phosphoric acid
DHxPA di-n-hexyl phosphoric acid
DHxOEPA di(hexoxyethyl) phosphoric acid
DOPA di-n-octyl phosphoric acid
~EPA mono-n-butyl phosphoric acid
MEHPA mono-2-ethylhexyl phosphoric acid
MOPA mono-n-octyl phosphoric acid
'i.,P;PA mono(p-1,l,3,3-tetramethylbutyl) phenyl
phosphoric acid
EHEH'PA 2-ethylhexyl 2-ethylhexyl phosphonic acid
DB'PA di-n-butyl phosphinic acid
TBP tri-n-butyl phosphate
TBPO tri-n-butyl phosphine oxide
THxPO tri-n-hexyl phosphine oxide
TOPO tri-n-octyl phosphine oxide
MiBC methyl isobutyl carbinol
MiBK methyl isobutyl ketone

the equilibrium conditions of the process. There are two general types

of approaches to studying the kinetics and equilibria of extraction.

The two types might be distinguished as "chemical" and "analog" or as

"theoretical" and "engineering." The "chemical" approach seeks rate

controlling steps as functions of the extraction variables, and aims

toward an actual chemical model of the extraction. The "analog"

approach seeks an adequate flow model of the extraction, evaluating

overall mass transfer coefficients for a material flux driven by concen-

tration gradients through resistants and evaluating phase diagrams and

operating curves at equilibrium in order to design the process. Both

approaches will be reviewed in the next two sections. Only information

pertinent to metal extraction by alkyl phosphoric acids has been included.

The reader is referred to Coleman and Roddy [35] and Treybal [36] for

more detailed discussions.

3-4 Theoretical Approach to Metal Extraction

3-4-1 Equilibria of metal extraction

.-~tn defining the equilibrium of mass transfer with chemical reaction,

it is necessary to formulate the appropriate equilibrium equations, and

find the associated equilibrium constant. In the following reactions,

HA represents organophosphoric acid, (org) and (aq) represent the organic

and aqueous phases respectively, C denotes concentration, and M denotes

the metal to be extracted. Several equilibria are involved

HA(aq) = H + A K = acid dissociation (3-3)

nHA(aq) = (HA)n(aq)

K = aqueous polymerization


nHa(org) = (HA) (org) Kn(org) = organic polymerization (3-5)

HA(aq) = HA(org) Kd = acid distribution (3-6)

Mm+ + (HA) (org) = M(AH nA n) (org) + mH+ Kr = organic reaction (3-7)
n n-i n-1 m r

M+ + (HA)n(aq) + M(AHnAnl)m(aq) + mH+ Kr(a) = aqueous reaction (3-8)

Reactions (3-7) and (3-8) can be combined into a distribution or

extraction coefficient

C (org)
EO = C- (3-9)
a CM(aq)

Kolarik [35] reviewed the studies which have been conducted to define

these equilibrium constants.

In practice the equations are rearranged and the K's combined to

give an extraction constant, K

[H ]m E nm
K = a (3-10)
e ((HA) mnCM(org))

Where [H ] and C (org) are equilibrium concentrations, and (HA) is
M i
the initial formal concentration of alkyl phosphoric acid, n is the

degree of polymerization of the organophosphoric acid, and m is the

charge of the metal. Similar expressions can be developed for monoalkyl

phosphoric acids. Extractions at low organic loading generally have a

polymerization value of 2 (n=2). This indicates that the organophosphoric

acid exists as a dimer, with one of the hydrogen ions available to react

with the metal. At higher metal to ester ratios the monomer to dimer

equilibrium is shifted to the left, and n has a value of 1. Baes et al. [37]

found that when the ratio of uranium to di(2-)thylhe:-y]) pho:iphoric acid

exceed 0.25 the alkyl phosphate reacted as a monomer rather than a dimer.

Baes and Baker [::] found that at a 0.167 ratio of iron (III) to

di(2-ethylhexyl) phosphoric acid the reaction changed from dimer to

monomer. They were able to exceed monomer loading (Fe/HiA = 0.33) and

extracted iron up to a ratio of 0.6, indicating that iron hydroxide

may have been extracted. As a rule, when the organic alkyl phos~p!lte

acts as a monomer rather than a dimer, the extractibility of the metal

is markedly lower [34].

3-4-2 Kinetics of metal extraction L-

The kinetic theory of simultaneous diffusion and chemical reaction

in the liquid phase was first developed by Hatta, Davis and Crandall,

and others (reported in [39]), based on the assumption that the resistance

to diffusion was concentrated within a film adjacent to the liquid-liquid

interface. The film was assumed to have negligible capacity for holding

the dissolved solute compared with the main body of the liquid, which

was so thoroughly mixed that no concentration gradient existed within it.

The assumption of two such films, one in each liquid,was developed by

Whitman. There are two bases for theoretical discussion of the kinetics,

one depending upon the assumption of physical films of negligible

capacity but finite resistance, and the other depending upon the assumption

of unsteady state molecular diffusion of the solute into the whole mass

of liquid. Since a physical picture completely consistent with the facts

is not available, neither theory can be completely accepted or refuted [39].

Much work has been done in developing the mathematics of simultaneous

mass transfer and chemical reaction based on both theories. Studies have

considered slow first order reversible and irreversible reactions,

slow nth order irreversible reactions, and infinitely rapid reversible

reactions [40]. Experiments have been conducted to test the validity

of some of these models [41,42].

Roddy et al. [43] studied the rate of iron extraction from a chemical

point of view. They defined several possible reactions that could take

place. By correlating extraction data to the predicted equations, they

were able to determine which reactions appeared to be rate controlling.

They concluded that the rate of extraction of iron (III) by di(2-ethylhexyl)

phosphoric acid was controlled by series and parallel steps at the

liquid-liquid interface during the introduction of the first and second

anion ligands in the formation of FeA3*3HA. The quantitative contri-

bution of each step to the net rate was given by a linear log-log equation.

A zero-order step with respect to iron was important with the quiescent

interface, but not with dispersion mixing. This implied that a build-up

of iron and mono-A complex at the quiescent interface limited iron

extraction, but this was not a limiting factor during dispersed mixing.

Kletenik and Navrotskaya [44], cited by Coleman and Roddy [35],

established that the rate was controlled by a chemical reaction at the

interface. They specifically examined and ruled out the possibility

of an iron-HA reaction in the bulk aqueous phase. These were the first

papers to propose an interface reaction, although the exact mechanism

was not understood.

3-4-3 Mechanism of metal extraction

The studies which imply that'the reaction takes place at the interface

have generally been ignored in the literature. The predominating theory

of mass transfer with simultaneous chemical reaction is based on the

two-film theory and has been described by several authors [34,39,45].

One of the basic assumptions of the two-film theory for the mechanism

of mass transfer is that the interface concentration of solute is at

equilibrium, i.e., there is no build-up of solute at the interface.

A second assumption is that the rate of mass transfer within each

phase is proportional to the difference in solute concentration in the

main body of the fluid and at the phase boundary or interface. The

transfer reaction is represented by equation (3-11).

M(aq) + X(org) = MX(orp) (3-11)

Figure 3-2 shows the mechanism of metal extraction. M is transferred

by molecular or eddy diffusion to the water phase film. By molecular

diffusion M moves through the film to the interface. X follows a

similar diffusion mechanism to reach the interface. The location at

which M and X meet may be either in the organic phase or the water phase,

depending upon which substance is diffusing slower. Figure 3-2 is drawn

such that the diffusion of X is rate controlling and the reaction zone

is therefore in the organic phase. Following chemical reaction (slow

or instantaneous) MX must diffuse back to the organic bulk. The process

can be M or X diffusion controlled or controlled by a chemical reaction.

However, there is no build-up of M at the interface. MX moves away at

a rate equal to or greater than the approach of M. New X must now

diffuse to the reaction zone to meet approach M.

A second approach to describing the mechanism of mass transfer with

chemical reaction is the surface-renewal theory which follows the same

assumptions as the two-film theory, except that M is brought all the

way to the interface by eddy diffusion. The solute is assumed to

Figure 3-2. Diagram of Metal Extraction as Proposed by the
Two-Film Theory.



instantaneously reach equilibrium and steady state. The film theory

and the surface-renewal theory cannot be distinguished by either the

form of the rate equations or the indicated effect of molecular

diffusivity [42]. The above mechanism is widely accepted, although

it has received some criticism. When the theoretical equations are

applied to data designed to test the validity of this mechanism,

invariably modifications must be made to the theory. The modifications

necessary for one set of data seldom agree with modifications made for

a second set of data. There seems to have been rationalization to

defend the model against the facts.

During the fifties (as outlined by Sherwood and Wei [46]) and

early sixties there was much evidence presented to contradict these

theories. And as noted in the last section, there was recent evidence

that the rate of metal extraction was controlled by a chemical reaction

rather than diffusion. Sherwood and Wei [46] stated that conclusions

seem to limit greatly the applicability of most of the existing theory

in cases of mass transfer between two liquids involving chemical

reaction. Modern text books still present the film theory as governing

the mass transfer of liquid-liquid extraction involving chemical reaction.

Until a new theory is shown to better explain the data, this model will

undoubtedly prevail.

3-5 Engineering Approach to Metal Extraction

3-5-1 Equilibria of metal extraction

In the case of two immiscible liquids, the equilibrium concen-

trations of a third component in each of the two phases are often related

by the extraction or distribution coefficient.

o 0
E (3-12)
a A

Where 0 = solute concentration in the organic phase and A = the solute

concentration in the aqueous phase. Equilibrium data are usually

reported in terms of the extraction coefficients. In the case of metal

extraction by chemical reaction, the coefficient is a function of several

variables, e.g., pH, concentration of solute, and concentration of alkyl

phosphoric acid. The coefficient should not be considered a constant.

The distribution coefficient does not apply when the two phases are

miscible. In many cases of practical importance the two phases are

partially miscible, and it is important to have solubility data as well

as information connecting the concentrations of an important component

in the two phases at equilibrium. The data are best expressed graphically.

Several graphical forms may be employed, but the most usual is the

ternary diagram.

In this research the choice of solvent was such that the two phases

can be considered completely immiscible, and the calculations are

therefore greatly simplified. Figure 3-3a shows an equilibrium curve

for two immiscible phases under batch extraction conditions. The curve

OP is the equilibrium curve which is found experimentally. Assume an

original solution of composition x contained in L liters of solvent A,

is reacted with G liters of pure solvent B in a single equilibrium contact.

The resulting mixture separates into two layers of equilibrium composition

x and yl. By a mass balance

L x = L x + G y (3-13)
o 1 1

Figure 3-3. Graphical Construction for Solvent Extraction
Design with Immiscible Liquids.

a. Crosscurrent multiple contact extraction

b. Countercurrent multiple contact extraction











1 L
xG (3-14)
x -Xo G

The composition of the original solution is represented by point R; the

composition xl and yl is obtained from the equilibrium curve OP by the

intersection of S of a line drawn through R with a slope L/G. Further

extraction of the raffinate T with pure solvent in a second equilibrium

contact will give raffinate and extract compositions x and y found
2 2
by the same method.

Figure 3-3b shows the equilibrium curve conditions for counter-

current extraction. By drawing a material balance around the first n

stages, it follows that

Yn+ = (Xn X ) + y (3-15)

The line representing this relation on a plot of y vs x is the operating

line and is based solely on the stoichiometry of the process. Figure 3-3b

represents the construction for a three stage countercurrent extraction

effecting the reduction of the raffinate concentration x to x3. The

line EF representing equation (3-15) is the operating line, drawn with

a slope L/G. Point E has the coordinates yn+, x where yn+ is the

solute concentration in the fresh solvent (usually zero) and xn is the

final raffinate concentration.

If it can be assumed that the two phases leaving each stage of

the contacting equipment are essentially in equilibrium, then the number

of equilibrium stages will be equal to the number of actual stages.

If equilibrium cannot be assumed, because of inadequate contact cf the

phases, then a "stage efficiency" may be introduced. This can be defined

as the ratio of the calculated number of stages to the actual stages


3-5-2 Kinetics of metal extraction

Generally, the rate of solvent extraction is reported only in

overall terms, such as stage efficiencies. It is possible to derive

an overall rate constant based on the two-film theory assumptions

presented in section 3-4-2. The equation is

JA = K A(C C *) = K A(C C ) (3-16)
A o o o w w w

Where JA is the rate of transfer as moles per unit time, A is the area

through which the solute transfers, K and K are the organic and water
o w

transfer coefficients, respectively, C is the bulk solute concentration

in the organic phase and C is the bulk solute concentration in the

water phase, C is the solute concentration of a solution in equilibrium
with C and C is the solute concentration of a solution in equilibrium
w w

with Co. The term JA/A is found experimentally, allowing the calculation

of K and K Although this equation is a purely imaginary picture of
o w

the process, it is possible to calculate useful transfer rates. Since

diffusion and first order chemical reactions have similar kinetics [41],

a slow first order chemical step will be equivalent to diffusion through

a fictitious film. In first order chemical systems therefore, the above

transfer coefficient may be used to express data provided that the

limitations of the constant are remembered. In general, however, the

overall mass transfer constant should be used only where transport to

and from the interface is the controlling factor, and not a chemical reaction.

For the purposes of calculating rate constants, the chemical

reaction is often assumed to be first order. Ryon et al. [47] derived

the following expression for batch systems

R 1
K = -ln(l-E) (3-17)


K = the effective rate constant,

R/V = the ratio of water to total volume,

t = time,


E = the stage efficiency defined as

C -C
E = (3-18)
C -C
o eq


C = initial aqueous solute concentration,

C = the equilibrium aqueous solute concentration,


C = the aqueous solute concentration at time t.

In Ryon's work the rate constant was directly proportional to the slope

of the curve obtained by plotting In(l-E) vs t. Batch rate constants,

although usually greater than those for continuous flow, are useful for

determining the relative effects of variables and to determine scale-up

factors. As Wells et al. [48] pointed out, this approach does not apply

when the reaction is pH dependent, as the reaction is not first order


3-6 Pertinent Solvent Extraction Literature

Very little work has been published on the extraction of aluminum.

The majority of the research has been done on expensive, high atomic

weight, radioactive materials such as uranium. In this review a few

examples of metal extraction research are presented along with a

discussion of the full scale uranium extraction processes. Finally,

the available literature on aluminum extraction is discussed.

Alkyl phosphoric acids were found to be effective uranium extractors

in work done at Oak Ridge in 1949. Subsequent investigations of the use

of alkyl phosphoric acids as liquid-liquid extraction reagents, primarily

for uranium, have been reviewed by Blake et al. [49,50]. The early work

on extraction by alkyl phosphates dealt mainly with di(2-ethylhexyl)

phosphoric acid extraction of uranium (VI) from acidic sulfate solutions.

Dialkyl phosphoric acids were thought to have extracted uranium by the

ion-exchange reaction

UO(aq) + 2HX(org) = UO2X2(org) + 2H+(aq) (3-19)

The data by Stewart and Hicks [51] and Blake [49,50] generally supported

this reaction. However, these authors also reported data which were not

in agreement with this reaction. In 1958, Baes et al. [37] pioneered

the work considering dimerization of the organic acid. Baes found that

up to a certain uranium:phosphoric acid ratio the uranium was extracted by

a dimer of alkyl phosporic acid as represented by equation (3-20).

After reaching thisratio, the reaction followed that of equation (3-19).

UO2+(aq) + 2(HX)2(org) = UO2X H2(org) + 2H+(aq) (3-20)

Between the (H ) r r, ;, of 0.4 M to 2 M and phosphoric acid concentration

range of 0.05M to 1M an equilibrium K of (4.0 0.4) was found for

reaction (3-20).

The Dapex process for uranium extraction was put into operation

in Colorado in 1956 [33]. Using 0.1 M di(2-ethylhexyl) phosphoric acid

it was possible to extract 4.5 to 7 grams of U02 per liter of kerosene.

The salts of the dialkyl phosphoric acid tended to precipitate or form

a third phase. However, this was prevented by the addition of a neutral

compound such as tributyl phosphate. It was also found that the neutral

additives had a very positive synergistic effect on extraction. The

uranium was removed from the kerosene by contacting the kerosene with a

carbonate solution. The carbonate solution was concentrated to 50 to

65 grams of U02 per liter.

A second process, also in commercial operation in 1958 was developed

by Dow [52]. Mono(2,6,8-trimethyl-4-nonyl) phosphoric acid at a

concentration of 0.1 M was used to extract the uranium. The kerosene

was loaded to 7 to 8 grams of UO2 per liter. Uranium was stripped from

the kerosene with 10 M hydrochloric acid. Sodium carbonate could not

be used for stripping because of the high solubility of the sodium

alkyl phosphate salt in the aqueous phase.

Wells et al. [48] tested several alkyl phosphoric acids for potential
use as selective extractants for beryllium from liquors high in aluminum

oxide. They concluded that the most suitable extractant was di(2-ethylhexyl)

phosphoric acid. Third phase formation was prevented by the addition

of 4% w/v octan-2-ol in the solvent. This modifier also enhanced

beryllium selectivity by repressing aluminum extraction. Decreasing

the pH of the ;queous phase also lowered aluminum extraction. Stripping

could be accomplished with either sodium hvydrioxie or ammou~ ium bifliurride.

Tests were conducted in a 3-in. diameter turbine at about 1800 rpm.

Detention time in the mixer was about 20 minutes. With an initial

concentration of 7.9 g BeO, 11 countercurrent stages could recover

96% of the BeO at a pH of 1. It was concluded that the extraction of

beryllium from sulfate liquors was practical.

Kolarik and Pankova [53] studied the extraction of lanthanides

by several dialkyl phosphoric acids. The mechanism of the extraction

was concluded to be

Ln3+(aq) + 3H A2(org) = Ln(HA2 )(org) + 3H (aq) (3-21)
22. 23

The log equilibrium K was found to be -1.15 to 4.76 for various lan-

thanides in a nitrate aqueous media. They found that increased branching

of the alkyl groups led to a substantial decrease in extraction effi-

ciency. Other authors have also found this result [54]. It had

previously been thought [34] that the lower efficiency with higher

branching was due to a difference in acidity of the phosphoric acid

compounds. However as Kolarik and Pankova [53] stated, the acidity

difference was very small, and cannot account for the drastic drop in

extraction efficiency. The extraction efficiency was not affected by

an increase in the number of carbon atoms in the alkyl groups of the same

structure. The extractability of metal was found to increase with

decreased radius of the extracted ions.

In 1968 George et al. [55] conducted a laboratory scale study of

the recovery of aluminum from copper mine acid waters. Aluminum was

extracted by a 0.4 N solution of monododecyl phosphoric acid. The optimum

extraction pH was 3.0 to 3.1. A kerosene:aqueous flow rate of 2:1 was

used. It was necessary to convert the phosphoric acid to the calcium

salt by addition of hydrated lime to the kerosene. This was to prevent

a pH drop during extraction which would lower the extraction coefficient.

With an initial aluminum concentration of 3.1 g/Z, the extraction

coefficient was 3.9. Two extraction stages were utilized, and one

liming stage. The stripping circuit was operated with three and four

stages and with 6 to 8 N HC1. Three stages proved adequate when the

organic to aqueous flow ratio was between 15:1 and 18:1. Under these

conditions, with a 10 minute contact time per stage, about 90% of the

aluminum was stripped by 6 N HC1. Stripping was not improved using

8 N HC1. The stripped kerosene was sent to a scrubbing stage before

returning to the extraction circuit.

Brown et al. [56] studied the rates of aluminum stripping as they

applied to aluminum's interference in the Dapex process for vanadium

recovery. In the process some aluminum was coextracted with uranium,

but the aluminum would not strip out during the stripping cycle and

therefore built-up as an impurity in the kerosene. In the Dapex process

for vanadium recovery, vanadium is extracted with 0.2 to 0.4 M

di(2-ethylhexyl) phosphoric acid and stripped with 1 M H2SO4 at ambient

temperature. At these stripping conditions, some of the impurity metals,

such as iron (III), aluminum and titanium are not removed effectively

from the organic phase. Tests were made to study the effects of temperature

on the stripping of these metals with 1 M H2SO Organic extractants

loaded with 0.8 g Al/liter were stripped at 30 C and 50 oC with an

organic:aqueous ratio of 1:1. Aluminum was about 25% stripped in

10 minutes at 30 oC, 60-75% in 20 minutes, and 85-95% in 60 minutes.


At 50 C, greater than 95% of the aluminum was stripped in less than

10 minutes. The rate of stripping increased at 30 C when tributyl

phosphate was added to the solvent phase.



The extraction experiments were divided into two categories:

aluminum extraction from synthetic feed solution, and aluminum ex-

traction from treatment plant sludges. All extractions were performed

in batch studies. Countercurrent extraction was simulated by the use

of batch extraction techniques. In the laboratory either of two

procedures to gather solvent extraction data could have been followed:

1) a batch simulation of the continuous multistage process could be

carried out; or 2) a continuous extraction in a miniature extractor

(bench-top scale) of known number of stages could be used. Both

procedures have been used in practice. However, miniature extractors

tend to be rather inflexible in the number of stages they have been

able to -represent and they have been shown to be difficult to feed

continuously in steady-state fashion on a very small scale [36].

Therefore, the batch simulation procedure was the method of choice for

this research. The batch simulation procedure called for repeated

introduction of feed mixture and solvents into a pattern of batch

extractions. These ultimately were able to produce the same effect as

a steady-state continuous process. The liquids in the batch process were

in the same volume ratio as the rates of flow to be simulated in a

continuous cascade. It has been shown that on a small scale there arc

ofLen relatively large losses of solution due to failure of the liquids

to completely drain from the vessels. Therefore, the volumes were

carefully monitored from stage to stage to insure that a constant

volume ratio was maintained. Figure 4-1 schematically shows the batch

simulation contact pattern that was followed in this research. The

batch simulation procedure asymptotically approached steady-state.

The number of horizontal rows required (in Figure 4-1) was dependent

upon each system studied.

4-1 Feed Solution

4-1-1 Synthetic feed solution

All synthetic feed solutions were made with distilled water and

reagent grade chemicals. All glassware was rinsed with tap water,

acid soaked in 50% HC1, washed with a cleaning solution, rinsed with

tap water, and rinsed with distilled water. Any glassware which had

contained organic compounds was acetone rinsed prior to the final

distilled water rinse.

Ten grams of aluminum metal wire were dissolved at elevated tem-

perature in 1 + 1 HC1 for preparation of the stock solution. In some

tests aluminum potassium sulfate was used as the source of aluminum.

The stock solution was prepared at 10,000 mg/P Al3+ (0.37 M). The
P04 ion was furnished to the stock solution from potassium phosphate

monobasic. The phosphate was added to give an Al:P weight ratio of

about 1.6:1. This was near the ratio to be expected for 90% phosphate

removal. The phosphorus concentration (as P) in the stock solution was

therefore 6200 mg/i-P (0.2 M). The pH of the stock solution was less

than 1.0, which ensured that aluminum phosphate precipitation did not

take place upon storage. When needed for experiments the stock solution

Figure 4-1. Batch Simulation of Three-Stage Continuous Counter-
current Cascade.

S = Solvent, F = Feed, R = Raffinate, E = Extract.



was diluted to the desired concentration with distilled water. pH

adjustments were made by sulfuric acid and/or sodium hydroxide addition.

In all extractions, unless otherwise stated, the Al:P weight ratio was

1.6:1. No further attempt to simulate ion concentrations in wastewater

was attempted, as actual treatment plant wastewater was to be used in the

subsequent set of experiments. The use of synthetic wastewater did allow

large quantities of feed solution to be readily available and the feed

parameters could be easily changed. Variability in feed quality could

be ignored when comparing different extraction parameters.

4-1-2 Tertiary wastewater sludge feed solution

The secondary effluent used for alum coagulation was collected

from the University of Florida Campus Sewage Treatment Plant. The plant

accomplished secondary treatment by both activated sludge (contact

stabilization) and trickling filter (standard and high rate). The sample

was collected from the plant's water reuse system. This water was

pumped from the final chlorine contact chamber. A typical analysis of

the secondary effluent is presented in Table 4-1. These values represent

a 10-month average, as reported in a previous study [57]. For this study,

the phosphate values were checked before and after each coagulation.

The secondary effluent was pumped into a 55-gallon drum for batch

coagulation. A 60 mg/1 alum dosage was applied to 40 gallons of effluent

per coagulation. The pH was not adjusted. Figure 4-2 shows alum

dosage vs final phosphate concentration for this effluent. The final

coagulation pH was approximately 6.00. The rapid mix time was 5 minutes,

flocculation time 15 minutes and the settling time was 2 hours. The

sludge was withdrawn such that the sludge volume was 1% of the coagulation

Figure 4-2.

Phosphate Removal as a Function of Alum Dose.
Solution is secondary effluent from the University
of Florida Campus Sewage Treatment Plant. Coagu-
lation pH = 6.0. Alum contained 8.0% aluminum




_. .

I 1.0-

0 30 .60 90 120


Characteristics of Secondary Effluent of the University of Florida
Campus Sewage Treatment Plant

Suspended solids



Kj eldahl-N




Total phosphate-P



20 mg/l

4.2 mg/Z

19.7 mg/k

5.6 mg/l

3.9 mg/l

6.5 mg/l

3.2 .. /

4.0 r:!Z

70 mg!/

31 mg/1

Source: Cornwell, D. A., Taylor, J. S., Furman,
T. deS., Zoltek, J., "Nutrient Removal
by Water Hyacinths," unpublished report,
Dept. Env. Eng., Univ. of Florida, 1973,
p. 25.

volume, or about 1.5 liters were withdrawn per batch coagulation. The

sludge was then returned to the lab for analysis. Acidification of

the sludge was always accomplished with 24 hours of collection, and

usually immediately after collection. When necessary, the sludge was

stored in a 3 C refrigerator.

4-2 Alkyl Phosphoric Acids

Five different alkyl phosphoric acids were supplied by Stauffer

Chemicals, Eastern Research Center, Westport, Connecticut. Properties

of the five acids are shown in Table 4-2 and 4-3. The di(2-ethylhexyl)

phosphoric acid that was actually used in the tests was from K and K

Chemicals, Plainview, New York. The properties of this acid as supplied

from K and K were not available, but should be close to those given in

Table 4-2 for Stauffer's Chemicals. All of these chemicals were

commercially available in large quantities at the time of this research.

The acids were not purified in any manner, but used directly as supplied

by the manufacturer. Since the chemicals would not be further purified

in full-scale application better design data could be obtained by using

the esters as supplied. The alkyl phosphoric acid solutions were

prepared by weighing out the appropriate amount of solution and diluting

it in kerosene. The kerosene was used as supplied from a local fuel

wholesaler. All molar solutions of alkyl phosphates were reported as

formal weights as given by the average molecular weights in Table 4-2.

The tributyl phosphate, [CH3(CH2)30]3P(O), was obtained from Aldrich

Chemicals, Atlanta, Georgia, and was reported as greater than 97% purity.

0 < 0 '0
C- 0 LP <
poo i a a
0 < 0 r D i 1u

0 00 CM It C0 C4
0 CM 00 M r l rH C-
4-1 4 r-I -T

00 .

o 4

0 cC 4J

10 0 *N)N 0 C ON 0

(J) 4- Z r

a aaC

Ss -H m 4


B 4- w ( -) O-

L O = 4-J 0 1 C, 4
rO fl 0O J r -a0 rC 0
SI o I rd -H -
1 01 MI H
O U 1U

*rl H R

r'U LO LA Cd

**H WJ -H a -
p- -i- i3 3 cd3 3 3

cnU U r']U cr cr -0 *iC)

r L C 4 QJ *J-l .- *-l H *rJ -
0 U 0 I 0_
1 01 Oa,

r-i o +-1
O I Oc I o H co CM o

TABL; 4-3

Solubility Properties of Alkyl

Acid Phosphates

Water Alcohol Acetone Ether CC14

Mono-di-iso-amyl I S S S S

Mono-di-2-ethylhexyl I S S S S

Mono-n-butyl S S S S SS

Mono-iso-amyl I S S S S

Di-2-ethylhexyl R SR S S S

Source: Stauffer Chemical Company, Akyl Acid Phosphates, Westport,
Conn., Aug., 1972.

Code: S = Soluble, SS = Slightly Soluble, R = Reacts, I = Insoluble.

4-3 Analytical Equipment and Techniques

4-3-1 Total phosphate determination

Total phosphate was determined by using acid perchlorate digestion,

followed by the vanadomolybdophosphoric acid colorimetric technique as

described in the thirteenth edition of Standard Methods for the Examina-

tion of Water and Wastewater. The sample was autoclaved for 30 minutes

at 15 psig and 120C in a "Castle thermatic 60" autoclave. The pH was

adjusted to just under the phenolphthalein endpoint. A sample (or diluted

sample) of 50 ml was reacted in a 50 ml graduated cylinder with 10 ml of

vanadatemolybdate for 10 minutes. The absorbance was read on a "Bausch

and Lomb Spectronic 70" at 400 nm. A light path of 2 cm was used.

4-3-2 Orthophosphate determination

Orthophosphates were determined exactly as total phosphates without

the inclusion of acid digestion.

4-3-3 pH measurement

pH was measured with a "Corning Model 12" expanded scale pH meter.

The pH meter was standardized at a pH of 2.00 for low pH determinations

and a pH of 7.00 for mid-range measurements. Rather than make an ionic

strength correction to determine the hydrogen ion concentration, the

proton activity as calculated from the measured pH was used in all cal-

culations. Therefore any equilibrium constant containing proton or

hydroxyl ion concentrations are reported as mixed constants.

4-3-4 Aluminum determination

The aluminum ion concentration was determined by atomic absorption

on a "Varian Techtron Model 1200." A nitrous oxide flame was used as

prescribed in the EPA manual, "!i.hiJds for Chemical Analysis of Water

and Wastes." All samples contained sufficient potassium ions to prevent

aluminum ionization. Readings were made at 309.2 nm. l-hen measuring

the aluminum in strong acid solution it was not possible to determine

the aluminum directly due to the interference from the high sulfuric

acid concentration (in the range of 1-10 normal). By diluting the sam-

ples one to one thousand with distilled water, accurate readings could

be achieved. The range of optimum aluminum determination was 10 to 150
mg/k as A13+

4-3-5 Available H determination

Available hydrogen ion determination of the alkyl phosphoric acid

was made by titration of the alkyl phosphoric acid-kerosene solution with

0.1 N NaOH. The sodium hydroxide was standardized by titration with

standard 0.1 N H2S04. The available hydrogen was that hydrogen that

would exchange with the sodium ion. The phosphoric acid acted as a dimer

such that equation (4-1) prevailed during titration.

NaOH(aq) + (HA)2(org) = NaAHA(org) + H20(aq) (4-1)

Millivolt and pH measurements were made using the "Corning Model 12"

expanded scale pH meter. Millivolt measurements could be made through-

out the titration. Aqueous phase pH measurements could only be made

after titration had proceeded far enough that sufficient water was

available to obtain an accurate reading. The point at which the alkyl

phosphoric acid had completed reaction with sodium was determined by

standard evaluation of acid-base titration curves [58]. In this pro-

cedure the inflection point of the titration curve represented the point

of complete acid neutralization.

4-3-6 Emulsion size determination

A "Unitron" microscope at a magnification of 1500 was used to meas-

ure droplet size. The microscope was set-up next to the mixing device.

While the organic-water solution was mixing, an eyedropper of solution

was taken from the beaker and a drop placed on the microscope slide.

The microscope was prefocused to allow rapid size determinations. It

was found that a more stable oil in water emulsion could be produced by

decreasing the oil:water ratio. The droplet size did not appear to be

affected by reducing the phase ratio, and more time was available to

obtain measurements. It was possible to obtain measurements down to

0.6 p at 1500 magnification.

4-3-7 Total solids determination

Total solids content of the sludge samples was determined by eva-

porating the water from the sludge at 1020C in an oven. In each case

a 50 gram (wet weight) sample was dried for 24 hours. The remaining

solids were weighed to determine milligrams of solids per 50 grams of

sludge. The percentage solids content could then be directly determined.

All solids contents, prior to thickening, were reported at a sludge

volume equal to 1% of the plant flow.

4-3-8 Volatile solids determination

The dried samples from section 4-3-7 were ignited at 5500C for

24 hours. The solids which remained after ignition represented non-

volatile solids. Volatile solids, which were equated to percentage

organic content of the sludge, were found by difference.

4-4 Experimental Procedure and Equipment

Batch extraction experiments were conducted in tall 300 ml beakers.

Samples of 50 ml volume were mixed with a variable speed mixer. Mixing

was at the maximum speed of the mixer which was about 1500 rpm, as re-

ported by the manufacturer. A polyethelyne three-bladed propeller was

used as the stirrer. The mixer was connected directly to a "Dimco-Gray"

timer to control mixing times. After mixing, the samples were immediately

transferred to 125 ml separatory funnels where the organic and water

phases were separated. A small volume of water was left in the funnel

after initial draining- The phases were then allowed to equalize before

draining the remaining water.

4-4-1 Treatment of wastewater sludge samples

Sludge samples were thickened using a "Phipps & Bird" jar stirrer

at 7 rpm. The supernatant was decanted using a siphon. The thickened

sludge was then acidified with 10 N H2SO4 to the desired pH. The sludge

was filtered through Whatman #2 filter paper. The sample was analyzed

for aluminum and phosphate content. The total available aluminum in the

sludge was determined by reacting 20 ml of sludge with 5 ml of concen-

trated HC1. The sample was filtered and analyzed for aluminum.

After filtration of the acidified sludge, some of the samples were

enhanced with additional aluminum and phosphate to simulate operation of

a treatment plant using a higher alum coagulation dose than was neces-

sary for proper coagulation of the Campus Sewage Treatment Plant effluent.

Enhancement was done by the addition of aluminum potassium sulfate and

potassium phosphate monobasic to the acidified filtered sludge.



Design of the recovery process required that the optimal operating

conditions for the extraction stages and the stripping stages be deter-

mined. In the extraction stages the aluminum was removed from the

acidified aluminum-phosphate sludge by dissolution into kerosene, re-

sulting in separation of the aluminum from the phosphate. In the

stripping stages the aluminum rich kerosene phase was contacted with

acid. The equilibria favored aluminum solubility in the acid, thereby

resulting in recovered alum. The aluminum free kerosene was recycled

back to the extraction stages.

Initial liquid extraction studies of aluminum were designed to

test the efficiency of various alkyl phosphoric acids as extractants.

In all aluminum extractions kerosene was the organic solvent of choice.

Kerosene has found wide use as a solvent for alkyl phosphoric acid dis-

solution. The alkyl phosphoric acids generally have high solubility in

kerosene; extraction coefficients are as high or higher with kerosene

than in other solvents; and kerosene has economic advantages over other


Both mono-n-butyl and mono-iso-amyl phosphoric acids have relatively

high water solubilities. It was found by Wells et al. [48] that the

dialkyl forms of these acids, dissolved in kerosene, would precipitate

with aluminum. Because of the loss of alkyl phosphate through precipi-

tation, the acids were undesirable for beryllium extraction. In initial

laboratory tests in this research, it was found that large amounts of

third phase formation developed when n-butyl and iso-amyl dihydrogen

phosphoric acids were used for aluminum extraction. .. 1ilf it may have

been possible to prevent third phase formation by a large amount of neu-

tral reagent addition, such as tributyl phosphate, these alkyl phosphates

were not further tested. The third acid mixture tested was an equal

molar mixture of mono- and di-iso-octyl phosphoric acid. It was found

that this acid mixture was not a good aluminum extractant in the hydrogen

form. A kerosene solution of 0.1 M iso-octyl phosphoric acid, converted
to the calcium salt, when contacted with 800 mg/Z (0.03 M) of Al at a

pH of 3, had an extraction coefficient of about 5.0. However, a 20% by

volume addition of tributyl phosphate (TBP) was necessary to prevent

third phase formation. Wells [48] also found that when di-iso-octyl

phosphoric acid extracted aluminum at a pH of 3, heavy precipitates

formed. At a pH of 1, he found light precipitates formed. Although

this acid could not be ruled out as an extractant for aluminum, other

alkyl phosphates tested proved to be more suitable for aluminum extrac-

tion. The next two acids tested both exhibited good extractant proper-

ties for aluminum. A section is devoted to each acid describing results

of extraction tests.

5-1 Aluminum Extraction by Di(2-ethylhexyl) Phosphoric Acid

Di(2-ethylhexyl) phosphoric acid (DEHFA) existed as a dimer in ker-

osene solution [37]. Therefore in its initial reaction with aluminum

the alkyl phosphate would act as a dimer. The monomer form of this acid

has been previously shown in Figure 3-la. A dimer form of the acid is

shown in Figure 5-la. In previous research, when DEHPA was used to ex-

tract Uranium (VI) [37] from solution, the reaction initially proceeded

Figure 5-1. Proposed Molecular Structure for a) DEHPA Dimer,
b) 1:4 U02 DEHPA Complex, and c) Polymeric
U02 DLHPA Complex. The anion X represents


The broken line represents hydrogen bonding.
Source: Baes, Jr., C. F., Zingaro, R. A. Coleman,
C. F., "The Extraction of Uranium (VI) from Acid
Perchlorate Solutions by Di(2-ethylhexyl) Phos-
phoric Acid in n-Hexane," J. Phy. Chem., 62:2,
134 (1958).



N. /
'< >*





U O2








with DEHPA acting as a dimer. Figure 5--b shows the proposed organic

complex resulting from (UO2)2+ extraction by DEHPA [37]. As the Uranium:

DEHPA ratio was increased the polymerzation of the Uranium-DL.'"PA complex

was believed to increase [37]. Figure 5-lc shows the proposed structure

of increased polymerzation. If aluminum extraction was similar to that

of UO 22+, the reaction could be represented by equation (5--1) at a low

Al:DEHPA ratio alkyll phosphoric acid is represented by HX throughout

this chapter).

Al(aq) + 3(HX)2(org).= AlX6H3(org) + 3H (aq) (5-1)

As the Al:DEHPA ratio surpassed that possible for a dimer reaction, the

DEHPA began to react as a monomer, as shown in equation (5-2).

Al3+(aq) + 3(HX)(org) = AiX3(org) + 3H+(aq) (5-2)

In several solvent extraction systems the reaction is pH dependent (for

example, uranium extraction is not pH dependent and beryllium extraction

is pH dependent [48]). pH control is especially important in pH depen-

dent systems where several countercurrent extractions are involved. If

the pH is allowed to drop during extraction then the pH of the feed solu-

tion to each stage will eventually fall below that for successful extrac-

tion. For example, in the beryllium extraction process [48] the initial

feed pH was 1.30, after four stages the pH had dropped to 0.61 which no

longer favored beryllium extraction. It was necessary to raise the pH

up to 1.0 before the feed entered stage 5. After stage 11 the pH had

dropped to 0.67, which was the last stage used for extraction. In this

research, rather than control aqueous feed pH between extraction stages,

the alkyl phosphoric acid was used in the salt form instead of the hydrogen

form. Since an acid stripping method was used the alkyl phosphoric acid

was in the hydrogen form and had to be converted to the salt form prior

to use as an extractant. This was done by direct addition of dry calcium

hydroxide to the organic solution. The alkyl phosphate reacted as a

dimer with the calcium such that,

2(HX)2(org) + Ca(OH)2 = CaX4H2(org) + 2H20 (5-3)

The subsequent aluminum extraction was,

3CaX4H2(org) + 2Al3+(aq) = 2AlX6H3(org) + 3Ca2+(aq) (5-4)

In the initial tests of aluminum extraction by DEHPA it was found

that extraction was a function of equilibrium pH. It was therefore nec-

essary to convert the DEHPA to the calcium form by following the sto-

ichiometry of equation (5-3). By adding Ca(OH)2(s) to the kerosene solution,

the maximum amount of calcium that could be reacted with the alkyl phos-

phoric acid was the amount indicated by equation (5-3). If more than a

1:4 Ca:DEHPA ratio was added, the amount of lime over this ratio acted

as excess base, raising the pH of the aqueous solution during extraction.

For a 0.1 M solution of DEHPA exactly 0.025 M Ca(OH)2 was needed to main-

tain a relatively constant pH during aluminum extraction. When calcium

di(2-ethylhexyl) phosphate (DEHPA-Ca) was formed in a kerosene solution

which was then contacted with a water phase the alkyl phosphate salt

tended to separate as a third liquid phase containing small amounts of

both water and kerosene. Third phase formation was prevented by the ad-

dition of a neutral organophosphorus compound, tributyl phosphate (TBP).

The amount of TBP needed was found to be directly proportional to the

concentration of DEHPA-Ca. The relationship between percent of TBP needed

on a volume/volume addition to kerosene as a function of the DEHPA con-

centration is shown in Figure 5-2. The aluminum extraction coefficient

obtained with 0.1 M DEHPA did not vary between 0% and 20% TBP addition.

No change in aluminum extraction was obtained between a mix time of

2 minutes and one hour. A 10-minute mix time for aluminum extraction by

DEHPA was used in the tests to represent equilibrium conditions. While

the values reported may not be true equilibrium values, they were con-

sidered to be close to equilibrium and relative changes in extraction by

changing various parameters can be evaluated by using the values deter-


The percentage of aluminum extracted as a function of equilibrium

pH is shown in Figure 5-3. The pH was controlled by varying the aqueous

pH and by varying the percentage of DEHPA that was converted to the salt

form. From this curve the importance of converting the DEHPA to the salt

form can be seen. At a raffinate pH of 1.6 only 33% of the aluminum was

extracted, while at a raffinate pH of 4.0 nearly 100% could be extracted.

It was not possible to attempt extraction above a pH of 4.0 due to

aluminum-hydroxide precipitation from the aqueous phase above this pH.

The extraction coefficient for a given aluminum concentration as a

function of DEHPA concentration was determined. Figure 5-4 shows extrac-

tion for a single contact as a function of DEHPA concentration at a feed

pH of 2.80. The initial aluminum concentration was 800 mg/ (0.03 M) for

each case in Figure 5-4 and the phosphate concentration was 500 mg/k

(0.016 M).

Many of the design data are presented at the feed pH rather than the

raffinate pH. It was felt that the feed pH is of more interest in deter-

mining design parameters. Even with the DEHPA in the salt form the

Figure 5-2. Percent of Tributyl
Prevent Third Phase
(DEHPA-CA) Complex.

Phosphate (TBP) Needed to
Formation of Calcium-DEAtPA
Aqueous pH = 2.83.


>20- o


- 15-



0 0.1 0.2 0.3 0.4

Figure 5-3.

Percent Aluminum Extracted by DEHPA as a Function
of pH. 0.1 M DEHPA, initial A13+ = 800 mg/j (0.03 M),
phase ratio = 1:1.


100 ,



1- 0

- 70 -


S50 o


1.5 2.0 2.5 3.0 3.5 4.0